IW 

Richard  M.  Holinan 


. 


A  COURSE  IN 

INORGANIC    CHEMISTRY 

FOR  COLLEGES 


BY 
LYMAN    C.     NEWELL,    PH.D.    (JOHNS  HOPKINS) 

-A 

PROFESSOR  OF  CHEMISTRY,  BOSTON  UNIVERSITY,  BOSTON, 
MASS.,  AUTHOR  OF  "  EXPERIMENTAL  CHEMISTRY  " 
u  DESCRIPTIVE    CHEMISTRY,"    "  GENERAL 
CHEMISTRY,"    "  LABORATORY   MAN- 
UAL OF  INORGANIC  CHEMISTRY 
FOR  COLLEGES" 


RE  VISED 


D.    C.    HEATH   &   COMPANY,  PUBLISHERS 
BOSTON  NEW  YORK  CHICAGO 


COPYRIGHT,  1909,  1916, 
BY  LYMAN  C.   NEWELL. 


PREFACE 

THIS  book  is  intended  for  college  students  who  devote  a 
year  to  general  chemistry.  It  is  primarily  a  students'  book, 
and  is  not  designed  to  replace  the  large  text-books,  which  are 
better  suited  for  reference  than  for  class  use.  The  descriptive 
portions  of  the  text  include  the  well-established  topics  usually 
taught  in  a  year  of  college  chemistry,  though  considerable 
space  is  devoted  to  the  application  of  modern  principles  to 
chemical  industries.  The  theoretical  portions  include  not 
only  the  principles  whose  value  was  demonstrated  long  ago, 
but  also  many  recent  conceptions,  which  are  fast  becoming 
indispensable  in  interpreting  chemical  and  physical  phe- 
nomena. Some  of  the  theoretical  topics,  which  are  distributed 
through  the  text  at  serviceable  points,  are  the  theory  of 
electrolytic  dissociation,  reversible  reactions,  equilibrium,  ca- 
talysis, vapor  pressure,  electrolysis,  and  the  behavior  of  dis- 
solved substances.  It  is  hoped  that  the  book  as  a  whole 
will  be  found  adequate  for  all  students  whose  work  is  con- 
fined to  a  year  and  will  likewise  serve  as  a  broad  foundation 
for  those  who  continue  the  study  of  chemistry. 

In  the  preparation  of  the  manuscript  and  correction  of  the 
proof,  judicious  advice  was  received  from  Professor  Frank  W. 
Durkee,  Tufts  College,  Tufts  College,  Mass.,  Assistant  Pro- 
fessor William  Foster,  Princeton  College,  Princeton,  N.J., 
Professor  Arthur  J.  Hopkins,  Amherst  College,  Amherst, 
Mass.,  Professor  James  F.  Norris,  Simmons  College,  Boston, 
Mass.,  Professor  Charlotte  F.  Roberts,  Wellesley  College, 


iv  PREFACE 

Wellesley,  Mass.,  Professor  Benjamin  W.  Van  Riper,  Wheaton 
College,  Wheaton,  111.,  and  others.  The  author  is  grateful  to 
these  teachers  for  their  help,  but  he  assumes  entire  responsi- 
bility for  all  statements  in  the  book. 

L.  C.  N. 

BOSTON  UNIVERSITY,  BOSTON,  MASS., 
January,  1909. 


PREFACE   TO    REVISED   EDITION 

OPPORTUNITY  has  been  taken  in  this  edition  to  revise  the 
portions  dealing  with  the  theory  of  chemistry.  Some  topics 
have  been  rewritten  and  several  new  ones  have  been  intro- 
duced. This  revision  and  extension  include  catalysis,  osmotic 
pressure,  hydrolysis,  colloidal  solutions,  mass  action,  reversible 
reactions,  displacement  of  equilibrium,  solubility  product,  ad- 
sorption, radioactivity,  atomic  weights,  valence,  and  molecular 
weights.  Many  other  topics  have  been  improved  and  extended 
to  conform  to  the  advance  of  science,  especially  those  dealing 
with  the  applications  of  chemistry  to  the  arts  and  industries 
as  well  as  to  life  itself.  Nearly  two  hundred  and  fifty  new 
problems  and  exercises  have  been  inserted.  Numerical  data 
have  been  corrected  and  the  tables  have  been  revised.  These 
extensive  modifications,  however,  have  not  destroyed  the  peda- 
gogical features  that  commended  the  first  edition  to  such  a 

large  number  of  college  teachers. 

L.  C.  N. 

BOSTON  UNIVERSITY,  BOSTON, 
May,  1916. 


CONTENTS 

CHAPTER  PAGE 

I.     MATTER,  ENERGY,  AND  CHANGE.        .  1 

Properties  of  Matter  —  Physical  and  Chemical  Changes 

—  Chemical  Action  —  Matter  and  Energy  —  Chemical 
Elements  —  Symbols  —  Chemical  Compounds. 

II.     OXYGEN  —  OZONE 15 

Preparation  and  Properties  of  Oxygen,  and  Chemical 
Changes  illustrated  by  Them  —  Oxides  and  Oxidation  — 
Combustion  —  Oxygen  and  Life  —  Uses  —  Discovery  — 
Ozone. 

III.  HYDROGEN    ....        .        .        .        .      30 

Preparation  and  Properties  of  Hydrogen,  and  Chemical 
Changes  illustrated  by  Them  —  Occlusion  —  Diffusion  — 
Oxyhydrogen  Blowpipe  —  Reduction. 

IV.  SOME  PROPERTIES  OF  GASES         .         .         .         .       41 

Properties  of  Gases  —  Gas  Volumes,  Temperature,  and 
Pressure  —  Standard  Conditions  —  Law  of  Charles  — 
Absolute  Zero  : —  Law  of  Boyle  —  Reduction  of  Gas 
Volumes  —  Weight  of  a  Liter  of  Oxygen  —  Densities  of 
Gases. 

V.     GENERAL  PROPERTIES  OF  WATER        ...      50 

Occurrence  and  Functions  in  Nature  —  Industrial  Appli- 
cations—  Natural,  River,  and  Drinking  Waters — Distil- 
lation —  Physical  Properties  —  Vapor  Tension  and 
Vapor  Pressure  —  Chemical  Properties — Dissociation  by 
Heat — Solvent  Power  —  Solution  of  Gases  —  Henry's 
Law  —  Solutions,  Liquids,  and  Solids  —  Crystallization 

—  Efflorescence  —  Solution   and  Vapor  Pressure  —  De- 
liquescence —  Therma)  Phenomena  —  Chemical  Action. 

v 


iv  PREFACE 

Wellesley,  Mass.,  Professor  Benjamin  W.  Van  Riper,  Wheaton 
College,  Wheaton,  111.,  and  others.  The  author  is  grateful  to 
these  teachers  for  their  help,  but  he  assumes  entire  responsi- 
bility for  all  statements  in  the  book. 

L.  C.  N. 

BOSTON  UNIVERSITY,  BOSTON,  MASS., 
January,  1909. 


PREFACE   TO    REVISED    EDITION 

OPPORTUNITY  has  been  taken  in  this  edition  to  revise  the 
portions  dealing  with  the  theory  of  chemistry.  Some  topics 
have  been  rewritten  and  several  new  ones  have  been  intro- 
duced. This  revision  and  extension  include  catalysis,  osmotic 
pressure,  hydrolysis,  colloidal  solutions,  mass  action,  reversible 
reactions,  displacement  of  equilibrium,  solubility  product,  ad- 
sorption, radioactivity,  atomic  weights,  valence,  and  molecular 
weights.  Many  other  topics  have  been  improved  and  extended 
to  conform  to  the  advance  of  science,  especially  those  dealing 
with  the  applications  of  chemistry  to  the  arts  and  industries 
as  well  as  to  life  itself.  Nearly  two  hundred  and  fifty  new 
problems  and  exercises  have  been  inserted.  Numerical  data 
have  been  corrected  and  the  tables  have  been  revised.  These 
extensive  modifications,  however,  have  not  destroyed  the  peda- 
gogical features  that  commended  the  first  edition  to  such  a 

large  number  of  college  teachers. 

L.  C.  N. 

BOSTON  UNIVERSITY,  BOSTON, 
May,  1916. 


CONTENTS 

IAPTEB  PAGE 

I.     MATTER,  ENERGY,  AND  CHANGE  .        .  1 

Properties  of  Matter  —  Physical  and  Chemical  Changes 

—  Chemical   Action  —  Matter  and  Energy  —  Chemical 
Elements  —  Symbols  —  Chemical  Compounds. 

II.     OXYGEN  —  OZONE          .        .     :  -...,.        .        .        .       15 

Preparation  and  Properties  of  Oxygen,  and  Chemical 
Changes  illustrated  by  Them  —  Oxides  and  Oxidation  — 
Combustion  —  Oxygen  and  Life  —  Uses  —  Discovery  — 
Ozone. 

III.  HYDROGEN .30 

Preparation  and  Properties  of  Hydrogen,  and  Chemical 
Changes  illustrated  by  Them  —  Occlusion  —  Diffusion  — 
Oxyhydrogen  Blowpipe  —  Reduction. 

IV.  SOME  PROPERTIES  OF  GASES         .         .         .         .       41 

Properties  of  Gases  — Gas  Volumes,  Temperature,  and 
Pressure  —  Standard  Conditions  —  Law  of  Charles  — 
Absolute  Zero  —  Law  of  Boyle  —  Reduction  of  Gas 
Volumes  —  Weight  of  a  Liter  of  Oxygen  —  Densities  of 
Gases. 

V.  GENERAL  PROPERTIES  OF  WATER  ...  50 
Occurrence  and  Functions  in  Nature  —  Industrial  Appli- 
cations—Natural,  River,  and  Drinking  Waters— Distil- 
lation —  Physical  Properties  —  Vapor  Tension  and 
Vapor  Pressure  —  Chemical  Properties — Dissociation  by 
Heat — Solvent  Power  —  Solution  of  Gases — Henry's 
Law  —  Solutions,  Liquids,  and  Solids  —  Crystallization 

—  Efflorescence  —  Solution   and  Vapor   Pressure  —  De- 
liquescence —  ThermaJ  Phenomena  —  Chemical  Action. 


vi  CONTENTS 

CHAPTER  PAGE 

VI.     COMPOSITION  OF  WATER  —  HYDROGEN  DIOXIDE       78 

Composition  —  Electrolytic  Decomposition — Volumetric 
and  Gravimetric  Composition  —  Morley's  Determination 

—  Hydrogen  Dioxide. 

VII.     LAW  AND  THEORY 88 

Law,  Theory,  Hypothesis  —  Law  of  Definite  Proportions 

—  Law   of   Multiple  Proportions — Atomic  Theory  — 
Atoms  and   Molecules  —  Atomic   Weights  —  Chemical 
Symbols  and  Formulas  —  Molecular  Weights  —  Calcula- 
tions—  Equations  —  Four  Kinds  of  Chemical  Change  — 
Making  Equations. 

VIII.     ATMOSPHERE  —  ARGON  —  NITROGEN    .        .        .     115 
Atmosphere  —  Properties  and  Ingredients  —  Properties 
of  Nitrogen  —  Oxygen  and  Nitrogen  in  the  Atmosphere 

—  Volumetric  and  Gravimetric  Composition  —  Water 
Vapor  and  Carbon  Dioxide  in  the  Atmosphere  —  Argon, 
Helium,  and  Related  Gases  —  Air  is  a  Mixture  —  Liquid 
Air — Liquefaction  of  Gases  —  Nitrogen,   Preparation, 
Properties,  and  Relation  to  Life — Fertilizers. 

IX.     SOLUTION  —  THEORY    OF    ELECTROLYTIC   DISSO- 
CIATION  132 

General  Properties  of  Solutions  —  Theory  of  Electro- 
lytic Dissociation  —  Ions  —  Electrolytes  and  Non-elec- 
trolytes —  Osmotic  Pressure  —  Freezing  Point  and 
Boiling  Point  —  Electrolysis  of  Solutions  —  Migration 
of  Ions  —  Chemical  Behavior  of  Electrolytic  Solutions — 
Common  Ions  —  Summary. 

X.     ACIDS,  BASES,  SALTS  —  NEUTRALIZATION     .        .    149 

Acids,  Properties,  Composition,  Definition  —  Bases, 
Properties,  Composition,  Definition  —  Salts,  Properties, 
Composition  —  Neutralization — Heat  of  Neutralization 

—  Classification  of  Salts  —  Basicity  and  Acidity  —  No- 
menclature of  Acids,  Bases,  and  Salts  —  Anhydrides  — 
Degree  of  Dissociation  — Salts  and  Electrolytic  Dissoci- 
ation —  Hydrolysis. 


CONTENTS  vii 

CHAPTER  PAGE 

XI.     ENERGY    AND    CHEMICAL    CHANGE  —  CHEMICAL 

EQUILIBRIUM 169 

Light  and  Chemical  Change  —  Heat  and  Chemical 
Change — Electric  Furnace  —  Measurement  of  Heat 
Energy  —  Exothermic  and  Endothennic  Compounds  — 
Electricity  and  Chemical  Change  —  Voltaic  and  Elec- 
trolytic Cells  —  Electrotyping  and  Electroplating  — 
Faraday's  Law  —  Catalysis  —  Equilibrium  —  Mass  Ac- 
tion —  Displacement  of  Equilibrium  —  Solubility 
Product. 

XII.     CHLORINE  AND  ITS  COMPOUNDS  ....     196 

Chlorine — Chlorine  Water — Liquid  Chlorine  —  Bleach- 
ing Powder  and  Bleaching  —  Hydrochloric  Acid  and 
Chlorides  —  Test  for  Ionic   Chlorine — Chlorine   Com- 
pounds. 

XIII.  COMPOUNDS  OF  NITROGEN  —  G-AY-LussAc's  LAW    210 

Ammonia  —  Liquid  Ammonia  —  Ammonium  Com- 
pounds —  Ammonium  Hydroxide  —  Refrigeration  — 
Composition  of  Ammonia  Gas  —  Nitric  Acid  and 
Nitrates  —  Nitrous  Acid  and  Nitrites  —  Oxides  of  Ni- 
trogen —  Aqua  Kegia  —  Gay-Lussac's  Law. 

XIV.  ATOMIC  AND  MOLECULAR  WEIGHTS  — VALENCE     227 

Atomic,  Molecular,  and  Equivalent  Weights — Kinetic 
Theory  of  Gases  —  Avogadro's  Hypothesis  —  Deter- 
mination of  Molecular  Weights  —  Victor  Meyer's 
Method  —  Freezing  Point  Method — Determination  of 
Atomic  Weights  —  Specific  Heat  — Molecular  Formulas 
—  Valence. 

XV.     CARBON  AND  ITS  OXIDES  —  CARBIDES         .         .     263 
Carbon  —  Diamond  —  Graphite  —  Coal  —  Charcoal  — 
Adsorption  —  Coke  —  Lampblack  —  Allotropism  — 
Carbon  Dioxide  —  Carbonic  Acid  —  Carbonates  —  Car- 
bon Monoxide  —  Calcium  Carbide. 

XVI.     HYDROCARBONS  —  ILLUMINATING  GAS  —  FLAME     288 
Methane  —  Ethylene  —  Acetylene,   Illumination,  Gen- 
eration —  Petroleum  —  Natural    Gas  —  Coal    Gas  — 
Water  Gas  —  Illuminating  Gases  —  Flame  —  Bunsen 
Burner  —  Oxidizing  and  Reducing  Flame. 


Vlll 


CONTENTS 


CHAPTER 

XVII. 


XVIII. 


XIX. 
XX. 

XXI. 
XXII. 

XXIII. 


PAGB 

OTHER  CARBON  COMPOUNDS     ....     310 

Organic  Compounds  —  Alcohols  —  Fermentation  — 
Formaldehyde  —  Ether — Acetic  Acid  and  Acetates 

—  Vinegar  —  Organic  Acids  —  Ethyl  Acetate  —  Fats, 
Glycerine,    and   Soap  —  Sugar  —  Starch  —  Bread  — 
Cellulose  —  Benzene  and  Related  Compounds  —  Cy- 
anogen —  Proteins. 

SULPHUR  AND  ITS  COMPOUNDS          .         .         .     332 

Sulphur,  Occurrence,  Source,  Extraction,  Purifica- 
tion, Properties  —  Forms  —  Transition  Point  —  Hy- 
drogen Sulphide  —  Sulphides  —  Sulphur  Dioxide  — 
Sulphur  Trioxide  —  Sulphurous  Acid  and  Sulphites 

—  Sulphuric    Acid  —  Chamber    Process  —  Contact 
Process  —  Sulphates  —  Sodium  Thiosulphate  —  Car- 
bon Bisulphide. 

CLASSIFICATION    OF    THE    ELEMENTS  —  PERI- 
ODIC TABLE,  GROUPS,  AND  LAW       .         .     355 

FLUORINE  —  BROMINE  —  IODINE       .        .        .     362 

Fluorine  —  Hydrofluoric  Acid — Bromine  —  Iodine  — 
Hydriodic  Acid  and  Equilibrium  —  Halogen  Family. 

BORON 374 

Boron  —  Boric  Acid  —  Borax  —  Borax  Beads. 

SILICON  —  GLASS 378 

Silicon  —  Silica  —  Silicic  Acids  and  Silicates  —  Col- 
loidal Solutions — Silicon  Tetrafluoride — Carborun- 
dum —  Glass. 


PHOSPHORUS  —  ARSENIC  —  ANTIMONY  —  BIS- 
MUTH       

Phosphorus  —  Oxides  —  Acids  and  Salts  —  Phos- 
phine  —  Phosphorus  Trichloride  —  Phosphorus  Pen- 
tachloride  and  its  Dissociation  —  Matches — Relation 
of  Phosphorus  to  Life  —  Arsenic  —  Oxide  —  Acids 
and  Salts  —  Sulphides  —  Marsh's  Test  —  Antimony 
and  its  Compounds  —  Bismuth  and  its  Compounds 
—  Nitrogen  Family. 


391 


CONTENTS 


ix 


CHAPTER 

XXIV. 
XXV. 


PAGE 

410 


METALS  AND  METALLURGY     . 

SODIUM  —  POTASSIUM  —  LITHIUM  —  AMMO- 
NIUM—  SPECTRUM  ANALYSIS  .  .  .  417 
Sodium  —  Chloride  —  Carbonate  —  Leblanc  and  Sol- 
vay  Processes  —  Bicarbonate,  Hydroxide,  Sulphate, 
Sulphite,  Nitrate,  Nitrite,  and  Peroxide  —  Potas- 
sium —  Stassfurt  Deposits  —  Potassium  Chloride, 
Nitrate,  and  Nitrite  —  Gunpowder  —  Chlorate,  Car- 
bonate, and  Hydroxide  —  Relation  to  Life  —  Lith- 
ium—  Ammonium  Chloride,  Sulphate,  Sulphide, 
Nitrate,  and  Carbonate  —  Spectrum  Analysis. 

XXVI.     COPPER  — SILVER  — GOLD      .         .        .        .442 
Copper —  Metallurgy  —  Purification  by  Electrolysis 

—  Cuprous  Compounds  —  Cupric  Sulphate  and  Ni- 
trate— Displacement  of  Metals  and  Electro-chemical 
Series  —  Silver  —  Metallurgy  —  Silver    Plating  — 
Silver  Nitrate  and  Hal  ides  —  Photography  —  Gold 

—  Mining  and  Metallurgy  —  Compounds  —  Copper 
Family. 

XXVII.     CALCIUM — STRONTIUM — BARIUM — RADIUM    463 
Calcium  —  Carbonate    and    Bicarbonate  —  Oxide 

—  Cement  —  Hydroxide  —  Sulphate  —  Chloride  — 
Hardness  of  Water  —  Strontium  —  Barium  —  Alka- 
line Earth  Family  —  Radium. 

XXVIII.     MAGNESIUM  —  ZINC  —  CADMIUM  —  MERCURY    478 
Magnesium  —  Oxide,  Hydroxide,    Sulphate,   Chlo- 
ride,  and   Carbonate  —  Zinc  —  Metallurgy  —  Com- 
pounds —  Cadmium  —  Mercury  —  Amalgams  — 
Compounds  —  Tests  —  Zinc  Family. 


XXIX.     ALUMINIUM     ....... 

Occurrence  —  Metallurgy  —  Oxide  —  Gems  —  Hy- 
droxide —  Hydrolysis  —  Aluminates  —  Sulphate  — 
Alum  —  Cryolite  —  Chloride  —  Clay,  Porcelain,  and 
Pottery. 


493 


CONTENTS 


CHAPTER  PAGE 

XXX.     TIN  AND  LEAD  —  CERIUM  AND  THORIUM     .     505 

Tin  —  Metallurgy  —  Transition  Point  —  Alloys  — 
Compounds — Oxidation  and  Reduction  —  Lead  — 
Metallurgy  —  Drinking  Water  —  Alloys  —  Oxides, 
Carbonate,  Sulphate  —  Cerium  and  Thorium. 

XXXI.     MANGANESE 517 

Preparation,  Properties,  Uses — Manganese  Dioxide 

—  Potassium  Permanganate  —  Compounds. 

XXXII.     CHROMIUM  —  URANIUM  —  RADIOACTIVITY     .     522 

Chromium,  Preparation,  Properties,  Uses  —  Potas- 
sium Chromate  and  Dichromate  —  Chrome  Alum 

—  Lead    Chromate  —  Compounds  —  Molybdenum, 
Tungsten,  Uranium  —  Radioactivity. 

XXXIII.  IRON  — NICKEL— COBALT      ....     532 

Occurrence  of  Iron  —  Metallurgy  —  Cast  Iron  — 
Wrought  Iron  —  Steel,  Processes,  Properties,  Uses  — 
Pure  Iron  —  Oxides,  Hydroxides,  Sulphates,  Sul- 
phides, Chlorides,  Carbonate,  and  Cyanides  — 
Nickel  —  Cobalt. 

XXXIV.  PLATINUM  AND  ASSOCIATED  METALS    .        .     554 

APPENDIX          .         . 557 

Metric  System — Thermometers — Crystallization  — 
Vapor  Pressure  —  Atomic  Weights. 


A  COURSE  IN 

INORGANIC  CHEMISTRY 

FOR  COLLEGES 


INORGANIC   CHEMISTRY 


CHAPTER  I 

Physical  and  Chemical  Changes  —  Matter  —  Energy  —  Ele- 
ments —  Compounds 

CHEMISTRY  is  one  of  the  natural  sciences.  By  a  natural 
science  we  mean  an  organized  group  of  knowledge  devoted 
primarily  to  the  study  of  matter  and  energy.  There  are 
several  of  these  groups.  Some,  like  geology,  deal  for  the 
most  part  with  the  concrete  aspects  of  science,  while  others, 
like  chemistry  and  physics,  are  concerned  with  both  abstract 
and  concrete  phases.  The  different  groups  of  natural 
sciences  are  not  independent.  Indeed,  they  often  overlap 
and  have  indefinite  boundaries.  This  has  always  been 
characteristic  of  chemistry  and  physics,  and  recent  study 
has  drawn  them  even  more  closely  together. 

The  scope  of  chemistry  is  rather  difficult  to  outline,  espe- 
cially at  the  outset  of  its  study.  Hence  it  is  only  possible 
at  this  stage  to  present  a  preliminary  and  somewhat  incom- 
plete conception,  leaving  to  future  pages  its  development, 
interpretation,  and  application.  Chemistry  deals  with  the 
properties  of  matter,  with  the  changes  involved  in  the  trans- 
formations of  different  kinds  of  matter,  with  numerous  laws, 
theories,  and  hypotheses  summarizing  chemical  phenomena, 
and  with  the  manufacture  and  utilization  of  a  vast  number  of 
substances  indispensable  to  mankind. 

Properties  of  Matter.  —  Different  substances  are  iden- 
tified by  characteristics  called  properties.  Thus,  some 

1 


INORGANIC   CHEMISTRY 

substances  are  solid,  others  liquid,  and  still  others  gaseous, 
though  the  special  physical  state  of  a  substance  depends 
usually  upon  its  temperature.  For  example,  water,  mercury, 
and  many  other  substances,  which  are  liquid  at  the  ordinary 
temperature,  become  solid  if  cooled  to  a  low  temperature 
and  gaseous  if  heated  to  a  high  temperature.  Other  familiar 
properties  are  color,  odor,  taste,  relative  weight,  crystalline 
structure,  luster,  hardness,  melting  point,  boiling  point,  and 
solubility.  Many  properties  are  exhibited  when  substances 
are  subjected  to  the  action  of  light,  electricity,  and  par- 
ticularly to  a  wide  range  of  temperature.  But  the  most 
important  properties  are  doubtless  those  exhibited  when 
different  substances  act  upon  each  other  and  thereby  produce 
the  profound  changes  characteristic  of  chemistry. 

Let  us  consider  some  properties  of  the  substance  sulphur. 
Observation  shows  it  is  a  solid  which  has  a  yellow  color,  no 
odor,  no  taste,  and  a  crystalline  structure.  It  is  relatively 
heavier  than  water,  because  it  sinks  when  placed  in  a  vessel 
of  water,  but  it  is  insoluble  in  water.  Experiment  shows 
that  when  sulphur  is  heated,  it  melts  into  a  pale  yellow 
liquid,  which  turns  brown  at  a  comparatively  low  tempera- 
ture and  remains  so  until  the  temperature  is  quite  high, 
whereupon  it  becomes  viscous  like  tar  and  finally  boils,  yield- 
ing a  yellow  smoke  which  looks  much  like  sulphur.  In  the 
light  sulphur  has  luster,  i.e.  it  reflects  light.  It  does  not 
conduct  electricity,  for  when  introduced  into  the  circuit  of 
an  ordinary  electric  bell,  it  prevents  the  electricity  from 
ringing  the  bell;  on  the  other  hand,  when  rubbed  with  a 
cloth,  sulphur  becomes  electrified  and  attracts  tiny  pieces  of 
paper.  If  a  piece  of  sulphur  is  heated  to  a  high  temperature, 
it  takes  fire  and  burns.  The  flame  is  blue,  and  an  invisible, 
suffocating  product  is  detected  by  the  odor;  if  burned  long 
enough,  all  the  sulphur  is  transformed  into  this  gas.  Finally, 
if  sulphur  and  a  powdered  metal,  such  as  iron  or  copper,  are 


PHYSICAL   AND   CHEMICAL   CHANGES  3 

mixed  and  heated  in  a  test  tube,  the  mixture  begins  to  glow, 
the  incandescence  often  spreading  throughout  the  mass 
even  after  the  test  tube  has  been  removed  from  the  flame. 
The  product  is  neither  sulphur  nor  iron  but  a  black  substance 
having  properties  quite  different  from  those  of  the  original 
constituents.  Thus,  step  by  step,  we  have  established  by 
observation  and  experiment  those  properties  which  serve 
to  identify  sulphur  and  to  distinguish  it  from  all  other  sub- 
stances. By  a  similar  though  often  more  elaborate  and 
complicated  procedure  the  properties  of  all  substances  can 
be  discovered,  recorded,  and  classified. 

Changes    in   Matter    and    Classification    of    Properties.  — 

Observation  and  experiment  show  that  substances  often 
change  or  can  be  changed.  These  changes  are  shown  by  a 
change  in  properties.  Sometimes  the  change  merely  in- 
volves a  temporary  change  in  properties.  Thus,  water  can 
be  changed  into  steam  or  ice,  and  iron  can  be  magnetized  or 
melted,  but  the  water  and  iron  are  not  thereby  fundamentally 
altered.  Apparently  they  are  changed  into  new  substances, 
but  they  are  not  essentially  changed,  for,  although  some  of 
the  properties  are  new,  rigid  examination  shows  that  no  new 
substances  have  been  formed.  The  steam,  ice,  magnetized 
iron,  and  melted  iron  are  still  the  substances  chemically 
known  as  water  and  iron.  Indeed,  it  is  only  necessary  to 
cool  the  steam  and  the  melted  iron  and  to  melt  the  ice  to 
obtain  the  original  substances,  which  are  familiar  as  water 
and  iron.  Changes  which  do  not  involve  the  transformation 
of  substances  into  new  substances  are  physical  changes. 
In  physical  changes  the  change  in  a  substance  is  often  only 
temporary  and  is  due  to  some  change  in  or  departure  from 
the  conditions  which  usually  prevail,  e.g.  a  change  in  tem- 
perature, pressure,  or  electrical  conditions,  and  as  soon  as 
these  new  conditions  are  removed  the  substance  regains  its 


4  INORGANIC   CHEMISTRY 

familiar  properties.  Thus,  when  iodine  is  heated  in  a  test 
tube,  the  solid  iodine  turns  into  a  beautiful  violet  vapor  which 
solidifies  in  the  colder  part  of  the  test  tube,  where  it  may  be 
recognized  as  iodine  by  its  steel-gray  color  and  crystalline 
form.  The  iodine  has  not  been  transformed  into  a  new  sub- 
stance, but  is  .merely  changed  physically,  assuming  for  the 
time  being  properties  characteristic  of  iodine  at  a  higher 
temperature  than  the  ordinary. 

Very  often,  however,  substances  are  changed  fundamen- 
tally, their  essential  nature  is  affected,  and  the  change  is 
manifested  by  the  formation  of  new  substances.  The  orig- 
inal substance  with  its  properties  disappears,  often  com- 
pletely, and  one  or  more  new  substances  with  characteristic 
properties  appear.  Thus,  coal  is  readily  changed  into  ash 
and  an  invisible  gas  which  have  properties  totally  unlike 
those  of  coal.  The  change  is  permanent,  too,  for  the  ash 
and  gas  are  new  substances  and  do  not  become  coal  again  as 
soon  as  they  are  cold.  Changes  which  involve  the  trans- 
formation of  substances  into  new  substances  are  chemical 
changes. 

An  examination  of  the  properties  of  sulphur  recorded 
above  reveals  two  classes  which  are  more  or  less  distinct. 
Color,  odor,  taste,  relative  weight,  crystalline  structure, 
luster,  and  electrical  behavior  are  exhibited  by  unchanged 
or  unchanging  sulphur.  They  are  associated  with  physical 
changes,  and  their  manifestation  does  not  involve  a  dis- 
appearance of  the  substance  nor  its  transformation  into 
another  substance.  Thus,  sulphur  is  readily  recognized  as 
sulphur  during  its  immersion  in  water  or  its  electrification 
by  rubbing.  Properties  exhibited  during  physical  changes 
or  by  an  unchanging  substance  are  called  physical  properties. 
On  the  other  hand,  certain  properties  — not  always  readily 
detected  —  are  exhibited  only  by  chemical  changes.  Their 
manifestation  involves  a  fundamental  change  in  the  nature 


PHYSICAL   AND   CHEMICAL   CHANGES  5 

of  substances.  Thus,  sulphur  burns  in  air  and  unites 
chemically  with  iron,  thereby  forming  new  substances; 
oxygen,  as  we  shall  see  later,  readily  undergoes  chemical 
transformation,  uniting  with  many  substances;  further- 
more, many  substances  are  decomposed  into  other  substances 
by  heat  or  electricity.  Those  properties  —  not  essentially 
physical  —  manifested  by  substances  when  they  undergo 
chemical  changes  are  called  chemical  properties. 

It  is  clear  from  the  foregoing  paragraphs  that  typical 
physical  changes  are  characterized  by  the  alteration  of 
physical  properties,  while  chemical  changes  are  character- 
ized by  the  formation  of  one  or  more  new  substances.  Ex- 
amples of  familiar  physical  changes  are  changes  in  physical 
state  (illustrated  by  the  formation  of  ice  and  vapor  from 
water,  and  vice  versa),  production  of  the  colors  in  the  sky, 
magnetization  of  iron,  and  electrification  of  glass.  Familiar 
chemical  changes  are  the  rusting  of  iron,  burning  of  oil  in  a 
lamp,  digestion  of  food,  souring  of  milk,  and  combustion  of 
wood. 

It  is  not  always  possible  to  separate  physical  and  chemical 
properties  into  two  sharply  defined  classes  nor  to  call  certain 
changes  entirely  physical  or  exclusively  chemical.  Some 
chemists  advocate  the  adoption  of  a  third  class  of  changes, 
viz.  physico-chemical  changes,  such  as  fusion  (i.e.  melting) 
and  solution.  Thus,  when  sulphur  is  boiled  or  dissolved  in 
some  liquid,  properties  are  exhibited  which  are  in  part  physi- 
cal and  in  part  chemical.  For  example,  if  boiled  sulphur  is 
poured  while  still  hot  and  thick  into  water,  the  product  is 
a  brown,  plastic,  gummy  mass  totally  unlike  the  original 
sulphur  in  appearance  though  it  reverts  to  ordinary,  yellow 
sulphur  in  time;  similarly,  powdered  sulphur  disappears 
(i.e.  dissolves)  in  the  liquid  carbon  disulphide,  and  is  ap- 
parently converted  into  another  substance,  though  it  can  be 
readily  recovered,  often  as  beautiful  yellow  crystals,  by  evap- 


6  INORGANIC   CHEMISTRY 

orating  the  liquid.  Many  substances  besides  sulphur  when 
fused  (i.e.  melted)  or  dissolved  exhibit  properties  which  do 
not  readily  fall  into  either  of  the  classes  mentioned  above. 
Such  properties  are  often  called  physico-chemical  properties. 

Chemical  Action.  —  When  emphasis  is  laid  upon  the  chemi- 
cal change,  the  participating  substances  or  reagents  are  said 
to  undergo  chemical  action,  to  interact,  or  to  react.  This 
mutual  or  reciprocal  action  is  called  a  chemical  reaction. 
Usage  of  these  terms  is  not  uniform,  but  they  are  used  in 
this  book  in  the  sense  just  stated.  Thus,  when  copper  is 
put  into  dilute  nitric  acid,  the  evidence  of  physical  change 
in  the  acid  is  the  rise  in  temperature,  but  this  is  quite  sub- 
ordinate to  the  evidence  of  chemical  change,  viz.  the  gradual 
disappearance  of  the  copper,  the  liberation  of  a  brown  gas, 
and  the  production  of  a  new  substance  which  dissolves 
readily  in  the  acid  and  colors  it  blue.  Further  experiment 
confirms  these  observations,  for  when  some  of  the  blue  acid 
solution  is  evaporated,  a  blue  solid  is  left  which  dissolves 
in  water,  gives  off  the  brown  gas  when  heated,  and  also 
forms  a  solution  which  deposits  copper  upon  an  iron  nail. 
Chemical  action  has  clearly  taken  place  and  the  whole  phe- 
nomenon was  a  chemical  reaction. 

Reactions  are  often  complicated,  but  extensive  study  has 
shown  that  there  are  four  general  kinds,  —  combination, 
decomposition,  substitution,  and  double  decomposition. 
These  will  be  illustrated  and  discussed  in  appropriate  places. 

Matter  and  Energy.  —  All  changes  studied  in  chemistry 
involve  matter  and  energy. 

It  is  sufficiently  accurate  for  our  present  purpose  to  define 
matter  as  substance  revealed  to  our  senses  by  properties,  the 
fundamental  property  being  weight.  In  chemical  changes 
matter  is  transformed ;  that  is,  substances  are  so  changed 
that  new  substances  are  formed.  Such  changes,  as  we  have 


PHYSICAL    AND    CHEMICAL    CHANGES  7 

already  learned,  involve  a  substitution  of  new  properties  for 
old  ones.  One  property  of  substances,  however,  is  not 
changed,  viz.  the  weight.  Careful  and  extensive  study  shows 
that  the  total  weight  of  the  substances  participating  in  a 
chemical  change  is  apparently  unaltered.  This  means  that 
the  sum  of  the  weights  of  the  substances  which  enter  into 
a  chemical  change  equals  the  sum  of  the  weights  of  the  new 
substances  resulting  from  the  chemical  change.  This  feature 
of  chemical  change  is  sometimes  briefly  summed  up  by  saying, 
"  matter  is  indestructible."  This  rather  comprehensive  state- 
ment is  one  form  of  a  law  called  the  Law  of  the  Conservation 
of  Matter.  It  is  preferably  stated  thus:  — 

No  weight  is  lost  or  gained  in  a  chemical  change. 

Observation  shows  that  work  is  done  by  motion,  heat, 
light,  and  electricity.  Thus,  at  Niagara  Falls  the  water 
turns  wheels  which  are  connected  with  dynamos,  and  the 
electricity  generated  thereby  is  used  to  operate  street  cars, 
furnish  light,  and  produce  the  heat  which  is  used  in  certain 
chemical  industries.  The  term  energy  is  applied  to  what- 
ever produces  work  or  can  be  converted  into  work.  For 
example,  motion,  heat,  light,  and  electricity  are  familiar 
forms  of  energy,  and  are  usually  studied  in  physics.  In 
chemistry  we  study  these  forms  and  also  another  form,  called 
chemical  energy.  Energy,  like  matter,  is  transformable, 
yet  apparently  constant  in  amount.  Observation  of  the 
actual  transformations  of  the  different  forms  of  energy  is 
an  everyday  experience.  Coal  in  burning  gives  up  its  chemi- 
cal energy  in  the  form  of  heat,  and  the  heat  becomes  motion. 
In  an  electric  battery  chemical  energy  is  changed  into 
electricity,  which  in  turn  becomes  motion  in  a  bell,  or  heat 
in  a  cooking  utensil,  or  light  in  a  metal  wire,  or  even  chemi- 
cal energy  in  an  electroplating  apparatus.  All  chemical 
changes  involve  a  transformation  of  energy.  The  chemical 
energy  is  locked  up,  so  to  speak,  in  substances,  and  when 


8  INORGANIC   CHEMISTRY 

chemical  changes  occur  there  is  a  transformation  and  redis- 
tribution of  energy.  Thus,  the  heat  which  is  often  needed 
to  start  chemical  changes  really  becomes  transformed,  in 
part  at  least,  into  chemical  energy,  while  the  heat  produced 
by  chemical  changes  is  the  result  of  the  transformation  of 
some  of  the  chemical  energy  into  heat.  So  also  the  light 
which  is  essential  in  photography  is  transformed  into  chem- 
ical energy  and  locked  up  as  such  in  the  chemicals  on  the 
photographic  plate.  Food  contains  chemical  energy,  and 
when  food  is  digesting,  this  energy  is  being  transformed  into 
heat,  which  maintains  the  temperature  of  the  body.  But 
all  these  transformations  involve  no  loss  or  gain  in  the  total 
amount  of  energy  involved.  When  a  known  amount  of 
heat  is  put  into  a  substance  and  stored  up  as  chemical 
energy,  as  it  were,  that  heat  can  be  recovered  as  heat  or  as 
an  equivalent  amount  of  some  other  form  of  energy.  The 
amount  of  chemical  energy  is  always  altered  in  chemi- 
cal changes,  but  the  total  amount  of  energy  involved 
is  unchanged.  The  characteristics  of  energy  are  often 
summarized  in  a  brief  statement  called  the  Law  of  the 
Conservation  of  Energy,  thus  :  — 

Energy  can  be  transformed  without  loss,  but  cannot  be  (by 
any  known  means)  created  or  destroyed. 

Chemical  Elements  and  Chemical  Compounds.  —  In  chem- 
istry we  study  almost  exclusively  two  classes  of  substances, 
viz.  elements  and  compounds.  Mixtures,  especially  the  kind 
of  mixture  known  as  a  solution,  receive  some  attention. 
Chemical  compounds  consist  of  chemical  elements  united 
with  one  another.  There  are  thousands  of  compounds,  but 
only  about  eighty  elements. 

Chemical  Elements.  —  An  extensive  examination  of  many 
different  kinds  of  matter  has  shown  that  certain  kinds  can 
be  decomposed  at  will  into  two  or  more  substances  totally 


PHYSICAL   AND    CHEMICAL    CHANGES  9 

unlike  the  original  matter,  but  that  other  kinds  are  not 
reducible  by  any  means  at  present  under  our  control.  Water, 
for  example,  can  be  decomposed  into  the  gases,  hydrogen  and 
oxygen,  which  are  entirely  different  from  water.  But  neither 
hydrogen  or  oxygen  can  be  decomposed  by  any  known  process. 
Furthermore,  oxygen  and  hydrogen  cannot  be  transformed 
into  each  other  by  any  known  process.  Other  substances 
can  be  added  chemically  to  oxygen  or  hydrogen,  but  nothing 
can  be  taken  away  from  them  chemically,  nor  can  anything 
be  produced  from  them  without  increasing  their  weight. 
When  oxygen  or  hydrogen  undergoes  chemical  change,  there 
is  addition,  an  increase  in  weight,  a  transformation  into  a  new 
substance  having  new  properties,  and  especially  a  greater 
weight  than  the  original  amount  of  oxygen  or  hydrogen; 
these  chemical  changes  are  never  accompanied  by  loss  in 
weight.  Furthermore,  oxygen  and  hydrogen  have  char- 
acteristic properties;  some  are  much  alike,  while  others  are 
totally  different.  Oxygen  and  hydrogen  are  elements,  and 
all  other  kinds  of  matter  which  have  fundamental  character- 
istics typified  by  oxygen  and  hydrogen  are  likewise  elements. 
The  essential  properties  of  the  chemical  elements  will  be 
set  forth  as  we  proceed  with  our  study  of  the  elements  and 
their  compounds.  It  has  been  customary  for  many  years 
to  regard  elements  as  those  substances  which  cannot  be 
decomposed  into  simpler  substances  or  transformed  without 
loss  of  weight.  It  would  follow  from  this  traditional  defini- 
tion that  the  elements  are  the  primary  forms  of  matter,  so 
to  speak,  and  are  chiefly  characterized  by  stability.  This  is 
true  of  most  elements,  as  far  as  we  know.  But  our  inability 
to  decompose  elements  does  not  necessarily  mean  that  they 
are  immutable  under  all  conditions.  Indeed,  their  instability 
is  shown  by  the  spontaneous  decomposition  of  certain  ele- 
ments, especially  radium.  Nevertheless,  it  is  sufficient  for 
our  present  purpose  to  regard  the  chemical  elements  as 


10 


INORGANIC    CHEMISTRY 


the  fundamental  materials  from  which  compounds  are  formed 
and  to  which  compounds  are  finally  reduced.  Questions 
concerning  transmutation  can  be  postponed  until  more 
facts  are  available. 

Each  element  is  designated  by  a  symbol,  which  is  an  abbre- 
viation of  its  name.     The  following  is  an  alphabetical 

TABLE  OF  IMPORTANT  ELEMENTS  AND  THEIR  SYMBOLS 


NAME 

SYMBOL 

NAME 

SYMBOL 

Al 

Lead    \ 

Pb 

Sb 

Lithium 

Li 

Arssnic                 •          •     • 

As 

Magnesium                       . 

Ms 

Ba 

Mn 

Bi 

Hg 

B 

Nickel  

Ni 

Bromine                         •     » 

Br 

Nitrogen  

N 

(  'juliiiiuni 

Cd 

Oxygen 

o 

Calcium          •     .    •     •     . 

Ca 

P 

c 

Platinum  

Ft 

Chlorine                             . 

Cl 

Potassium 

K 

C  li  ro  in  ium 

Cr 

Silicon           . 

Si 

Cobalt   

Co 

Silver  

Ag 

Copper            •          •,     .    . 

Cu 

Na 

Fluorine 

F 

Strontium      

Sr 

Gold 

Au 

Sulphur 

s 

H 

Tin 

Sn 

Iodine 

I 

Zn 

Fe 

The  elements  are  not  uniformly  distributed  in  nature, 
either  in  abundance  or  mode  of  occurrence.  Our  knowledge 
of  the  relative  abundance  of  the  elements  is  based  on  a 
study  of  the  atmosphere,  the  ocean,  and  a  shell  of  the 
earth's  crust  about  ten  miles  deep.  The  approximate  dis- 
tribution of  matter  in  these  three  portions  of  the  globe  is 
seen  in  the  following 


PHYSICAL   AND    CHEMICAL   CHANGES 


11 


TABLE  OF  THE  DISTRIBUTION  OF  MATTER 

Atmosphere 03  per  cent 

Ocean 7.08  per  cent 

Shell  of  Earth's  Crust 92.89  per  cent 

The  atmosphere  contains  about  20  per  cent  of  oxygen,  79 
per  cent  of  nitrogen,  and  1  per  cent  of  argon.     The  distribu- 
tion of  the  elements  in  the  ocean  appears  in  the  following 
TABLE  OF  THE  APPROXIMATE  COMPOSITION  OF  THE  OCEAN 


ELEMENT 

PER  CENT 

ELEMENT 

PER  CENT 

85.79 

.09 

10.67 

.05 

Chlorine    

2.07 

Bromine  . 

.008 

Sodium-                        •     . 

1  14 

Carbon    . 

002 

Magnesium    

.14 

Other  Elements     .     .    . 

traces 

Besides  the  elements  which  are  combined  as  saline  matter, 
the  ocean  contains  dissolved  gases,  such  as  oxygen,  nitrogen, 
and  carbon  dioxide. 

The  proportion  of  the  elements  in  the  ten-mile  shell  of  the 
earth's  crust  is  shown  in  the  following 

TABLE  OF  THE  APPROXIMATE  COMPOSITION  OF  THE  EARTH'S  CRUST 


ELEMENT 

PER  CENT 

GRAPHIC  PROPORTION 

Oxygen  . 

47.07 

OQ  n« 

Aluminium  . 

7.90 

- 

Iron    .     .     . 

4.43 

_ 

Calcium  .     . 

3.44 

— 

Potassium   . 

2.45 

— 

Sodium  .     . 

2.43 

— 

Magnesium 

2.40 

— 

Remainder  . 

1.82 

*  ^ 

These  elements  are  chiefly  in  the  combined  state. 


12 


INORGANIC  CHEMISTRY 


The  chemical  elements  necessary  for  human  beings  are 
shown  in  the  following 

TABLE  OF  THE  AVERAGE  COMPOSITION  OF  THE  HUMAN  BODY 


ELEMENT 

PER  CENT 

ELEMENT 

PER  CENT 

ELEMENT 

PER  CENT 

Oxygen  .     .     . 
Carbon  .     .     . 

65.00 
18.00 

Phosphorus  . 
Potassium 

1.00 
0.35 

Magnesium    . 
Iron      .     .     . 

0.05 
0.004 

Hydrogen  .     . 
Nitrogen    . 
Calcium     .     . 

10.00 
3.00 
2.00 

Sulphur  .  . 
Sodium  .  . 
Chlorine  .  . 

0.25 
0.15 
0.15 

Iodine  .     .     . 
Fluorine    .     . 
Silicon  •     .     . 

trace 
trace 
trace 

The  following  is  a 

TABLE  OF  THE  LESS  COMMON  ELEMENTS  AND  THEIR  SYMBOLS 


NAME 

SYMBOL 

NAME 

SYMBOL 

Arsron 

A 

Ra 

Beryllium 

Be 

Rhodium  .          .... 

Rh 

Csesium           

Cs 

Rubidium     

Rb 

Cerium                          •     . 

Ce 

Ru 

Columbium 

Cb 

Sm 

Erbium      

Er 

Scandium     

Sc 

Europium            .              . 

Eu 

Se 

Gadolinium                        • 

Gd 

Tantalum      

Ta 

Gallium      

Ga 

Tellurium     

Te 

Germanium 

Ge 

Terbium  .          .... 

Tb 

Helium            •                    . 

He 

Thallium  

Tl 

Indium 

In 

Th 

Iridium 

Ir 

Thulium  

Tm 

Kr 

Ti 

La 

W 

Molybdenum  . 

Mo 

U 

Neodymium 

Nd 

V 

Neon 

Ne 

Xe 

Os 

Ytterbium    

Yb 

Pd 

Yt 

Pr 

Zirconium     .          ... 

Zr 

PHYSICAL   AND    CHEMICAL    CHANGES 


13 


The  symbols  of  the  chemical  elements,  as  may  be  seen  in 
the  foregoing  tables,  are  in  some  instances  the  first  letter 
of  the  common  name  of  the  element.  Thus,  O  is  the  symbol 
of  oxygen,  H  of  hydrogen,  N  of  nitrogen.  But  since  several 
elements  have  the  same  initial  letter,  the  symbol  of  some 
elements  contains  two  letters.  Thus,  C  represents  carbon, 
while  the  symbol  of  calcium  is  Ca,  of  chlorine  Cl,  of  chromium 
Cr,  and  of  copper  Cu.  The  symbols  of  several  elements, 
especially  the  metals  so  long  known,  are  derived  from  their 
Latin  names.  Thus,  we  have  a 

TABLE  OF  THE  CHEMICAL  ELEMENTS  WITH  LATIN  SYMBOLS 


ELEMENT 

LATIN  NAME 

SYMBOL 

ELEMENT 

LATIN  NAME 

SYMBOL 

Antimony    . 

Stibium 

Sb 

Mercury 

Hydrargyrum 

Hg 

Copper   .     . 

•Cuprum 

Cu 

Potassium  . 

Kalium 

K 

Gold  .     .     . 

Aurum 

Au 

Silver     .    . 

Argentum 

Ag 

Iron   .     .     . 

Ferrum 

Fe 

Sodium  .     . 

Natrium 

Na 

Lead  .     . 

Plumbum 

Pb 

Tin    ... 

Stannum 

Sn 

The  symbols  of  the  elements  always  begin  with  a  capital 
letter,  and  are  not  followed  by  a  period.  They  should  be 
learned  by  actual  use.  Their  full  significance  will  be  ex- 
plained in  later  chapters. 

Chemical  Compounds.  —  A  chemical  compound  is  a  sub- 
stance which  is  composed  of  two  or  more  chemical  elements. 
There  is  a  very  large  number  of  chemical  compounds.  The 
elements  which  make  up  a  chemical  compound  are  called  its 
constituents.  Chemical  compounds  differ  fundamentally 
from  elements  in  having  three  essential  characteristics. 
(1)  Their  constituents  are  not  mixed,  but  are  united  chemi- 
cally. That  is,  the 'chemical  energy  originally  possessed  by 
the  elements  that  make  up  compounds  has  so  operated  that 


14  INORGANIC  CHEMISTRY 

the  new  substance  (i.e.  the  compound)  cannot  be  separated 
into  its  constituents  except  by  the  application  of  some 
form  of  energy,  such  as  heat,  light,  electricity,  etc.  Thus, 
water  is  a  compound  of  the  elements  oxygen  and  hydrogen, 
which  cannot  be  separated  from  each  other  unless  water  is 
subjected  to  intense  heat  or  to  electricity.  (2)  In  any  given 
chemical  compound  the  constituents  are  always  the  same 
and  are  present  in  a  fixed  ratio  by  weight.  Thus,  sodium 
chloride  contains  39.34  per  cent  of  the  element  sodium  and 
60.65  per  cent  of  the  element  chlorine.  (3)  The  properties 
of  a  compound  differ  from  those  of  its  constituent  elements. 
Thus,  the  blue  solid  copper  sulphate  is  a  compound  of  three 
elements,  —  the  red  metal  copper,  the  yellow  non-metal  sul- 
phur, and  the  colorless  gas  oxygen. 

Chemical  compounds  must  not  be  confused  with  mixtures. 
The  components  of  a  mixture  may  vary  in  nature  and  in 
proportion  ;  they  are  also  not  combined  chemically  but  held 
together  loosely,  and  can  be  separated  by  some  mechanical 
operation,  as  filtering  or  sifting,  without  bringing  about  any 
chemical  change  by  the  application  of  energy.  A  mixture, 
too,  generally  has  properties  which  are  similar  to,  or  are  an 
average  of,  those  of  its  components. 

PROBLEMS 

1.  Solve  the  problems  on  the  metric  system  in  the  Appendix,  §  1. 

2.  Assuming  that  the  ten-mile  shell  of  the  earth's  crust  weighs 
900  trillion  metric  tons,   calculate  the  weight  of  each  abundant 
element  in  it.     (See  page  11.) 

3.  Calculate  the  weight  of  each  element  in  a  human  body  weigh- 
ing 70  kilograms.     (See  page  12.) 

4.  Calculate  the  number  of  grams  in  52  kilograms  of  typical 
ocean  water. 


CHAPTER  II 
Oxygen  —  Ozone 

OXYGEN  has  played  an  important  part  in  the  development 
of  chemistry,  and  is  an  appropriate  element  with  which  to 
begin  a  systematic  study  of  this  science. 

Occurrence.  —  Oxygen  is  the  most  abundant  and  widely 
distributed  of  the  elements,  and  occurs  both  free  and  com- 
bined. Mixed  with  nitrogen  and  a  few  other  gases,  it  forms 
nearly  21  per  cent  (by  volume)  of  the  atmosphere.  Com- 
bined with  hydrogen,  it  constitutes  eight  ninths  (by  weight) 
of  water;  combined  with  silicon  and  certain  metals,  it  makes 
up  nearly  half  of  the  earth's  crust;  while  compounds  of 
oxygen  with  carbon  and  hydrogen  form  a  large  part  of 
animal  and  vegetable  matter.  Thus,  the  human  body  con- 
tains about  65  per  cent  of  oxygen  (see  page  12),  and  starch, 
which  is  a  component  of  all  plants,  contains  50  per  cent. 

Preparation.  — Oxygen  can  be  obtained  from  its  compounds 
or  from  air.  It  was  first  prepared  by  decomposing  a  com- 
pound of  oxygen  and  mercury  called  mercuric  oxide.  When 
heated,  this  compound  decomposes  into  oxygen  and  mercury; 
the  oxygen  escapes  as  a  gas,  and  the  mercury  condenses  as 
globules  or  a  film  on  the  upper  part  of  the  glass  vessel  in 
which  the  experiment  may  be  conveniently  performed.  This 
experiment  is  historically  interesting,  because  it  was  first 
performed  by  Priestley,  the  discoverer  of  oxygen.  (See 
Discovery  of  Oxygen,  below.) 

15 


16 


INORGANIC   CHEMISTRY 


The  gas  is  often  prepared  by  decomposing  potassium 
chlorate  —  a  compound  of  oxygen,  chlorine,  and  potassium. 
If  heated  to  a  rather  high  temperature,  the  potassium 
chlorate  passes  through  a  series  of  changes  which  result 
finally  in  the  liberation  of  all  the  oxygen  and  the  formation 
of  a  compound  of  potassium  and  chlorine  called  potassium 
chloYide. 

Oxygen  is  most  conveniently  prepared  in  the  laboratory  by 
heating  a  mixture  of  potassium  chlorate  and  manganese 


FIG.  1.  —  Apparatus  for  preparing  and  collecting  oxygen  in  the  laboratory. 
The  oxygen  mixture  is  put  in  A  and  heated.  The  oxygen  gas  escapes 
through  B-C-D  into  the  pneumatic  trough  (#).  The  latter  is  filled  with 
water  until  the  shelf  is  covered  and  a  bottle  full  of  water  is  then  placed 
mouth  downward  on  the  shelf.  The  oxygen  bubbles  through  the  hole  in 
the  shelf  up  into  the  bottle  and  displaces  the  water. 

dioxide  in  a  glass  or  metal  vessel,  and  collecting  the  liberated 
oxygen  in  a  bottle  by  means  of  a  pneumatic  trough.  (See 
Fig.  1.)  The  gas  is  liberated  freely  from  this  mixture  at  a 
lower  temperature  than  when  either  compound  is  heated 
alone.  Large  quantities  of  oxygen  are  prepared  for  com- 
mercial purposes  by  heating  a  mixture  of  potassium  chlorate 
and  manganese  dioxide  in  an  iron  vessel. 


OXYGEN  —  OZONE  17 

Other  commercial  processes  are  used.  In  Erin's  process, 
which  is  operated  largely  in  England,  purified  air  is  forced 
by  a  pump  over  barium  oxide  heated  to  about  700  degrees 
C.,1  thereby  forming  barium  dioxide.  The  air  supply  is 
then  cut  off,  and  the  pressure  in  the  retorts  reduced  by 
reversing  the  pump.  This  operation  changes  the  barium 
dioxide  back  again  into  barium  oxide  and  oxygen.  The 
gas  is  drawn  off  into  a  reservoir.  The  process  is  then 
repeated. 

Oxygen  can  be  prepared  from  water.  When  an  electric 
current  is  passed  through  water  containing  sulphuric  acid 
or  sodium  hydroxide,  oxygen  (and  also  hydrogen)  is  liberated. 

Oxygen  can  also  be  separated  from  liquid  air.  (See  Liquid 
Air.)  By  allowing  liquid  air  to  evaporate  at  ordinary  tem- 
perature and  pressure,  the  nitrogen  escapes  from  the  liquid 
air  more  rapidly  than  the  oxygen,  leaving  finally  a  liquid 
which  is  nearly  pure  oxygen.  Unlimited  quantities  of 
oxygen  may  probably  thus  be  cheaply  prepared  from  the 
air. 

Oxygen  can  be  prepared  from  sodium  peroxide  (or  from 
oxone  —  a  special  form  of  this  compound)  by  allowing  it 
to  interact  with  water.  This  is  a  convenient  method  for 
preparing  a  small  quantity  of  the  gas. 

Chemical  Changes  illustrated  by  the  Preparation  of  Oxy- 
gen. —  Brief  reference  has  already  been  made  to  the  four 
main  kinds  of  chemical  change,  viz.  decomposition,  combi- 
nation, substitution,  and  double  decomposition.  The  first 
and  second  kinds  are  illustrated  by  some  of  the  methods  for 
the  preparation  of  oxygen.  When  mercuric  oxide  is  heated, 

1  C.  is  the  abbreviation  of  "centigrade,"  which  is  the  name  of  the  ther- 
mometer used  in  science.  This  thermometer  registers  100  degrees  at  the 
boiling  point  of  water,  and  0  degrees  at  the  freezing  point.  (See  Appendix, 
§2.) 


18  INORGANIC   CHEMISTRY 

it  decomposes  into  oxygen  and  mercury.  Energy  in  the  form 
of  heat  serves  to  start  and  maintain  the  chemical  change, 
and  if  heat  were  supplied  long  enough,  all  the  compound 
would  be  decomposed  into  the  two  elements  oxygen  and 
mercury.  A  chemical  change  whereby  a  compound  with  a 
definite  set  of  properties  decomposes  into  two  elements  each 
with  its  own  set  of  properties  is  the  simplest  illustration  of 
the  kind  of  chemical  change  called  decomposition.  This 
fact  might  be  compactly  expressed  by  an  equation,  thus:  — 

Mercuric  Oxide  =  Mercury  +  Oxygen. 

This  equation  also  emphasizes  the  fact  that  the  sum  of  the 
weights  of  the  mercury  and  oxygen  equals  the  weight  of 
the  mercuric  oxide  decomposed  —  a  fact  readily  proved  by 
weighing  the  three  substances.  When  potassium  chlorate 
is  heated  until  all  the  oxygen  is  liberated,  the  chemical 
change  is  also  a  decomposition.  In  this  case,  however,  the 
products  are  the  element  oxygen  and  the  compound  of 
potassium  and  chlorine  called  potassium  chloride.  The  heat 
is  not  sufficient  to  produce  the  complete  decomposition  of 
potassium  chlorate  into  its  three  elementary  constituents. 
That  is,  decomposition  stops  when  all  the  oxygen  has  been 
liberated.  These  facts  might  be  expressed  thus  :  — 

Potassium  Chlorate  =  Potassium  Chloride  4-  Oxygen. 

Decomposition  in  general,  then,  is  a  chemical  change  whereby 
a  compound  is  separated  chemically  into  two  or  more  new 
substances.  The  preparation  of  oxygen  from  barium  oxide 
and  the  oxygen  of  the  air  illustrates  both  combination  and 
decomposition.  In  the  first  stage,  when  barium  oxide  is 
heated  in  a  current  of  air,  the  compound,  barium  oxide, 
unites  chemically  with  the  element,  oxygen,  and  thereby 
produces  a  new  compound  (barium  dioxide).  A  chemical 
change  whereby  two  or  more  elements  or  compounds  unite 


OXYGEN  — OZONE  19 

to  form  a  new  compound  is  called  combination.  The  sec- 
ond stage  in  the  preparation  of  oxygen  by  the  method  under 
discussion  is  another  illustration  of  decomposition.  Thus, 
barium  dioxide  decomposes  into  barium  oxide  and  oxygen. 
The  two  chemical  changes  might  be  expressed  thus  :  — 

(1)  Barium  Oxide  -f  Oxygen  =  Barium  Dioxide. 

(2)  Barium  Dioxide  =  Barium  Oxide  +  Oxygen. 

The  equations  used  in  this  paragraph  to  summarize  the  facts 
connected  with  decomposition  and  combination  are  crude 
forms  of  the  chemical  equation.  Subsequently  it  will  be 
shown  that  a  definite  relation  exists  between  the  weights  of 
the  substances  used  and  produced  and  that  this  relation  can 
be  expressed  accurately  in  the  form  of  an  equation.  At  the 
present  stage  the  rather  cumbersome  equation  used  above 
is  convenient  because  it  emphasizes  two  facts  about  chemi- 
cal change;  viz.  (1)  chemical  change  involves  no  change  in 
the  total  weight  of  the  substances,  and  (2)  chemical  change 
results  in  the  formation  of  new  substances. 

Properties.  —  Pure  oxygen  gas  has  no  color,  odor,  or  taste. 
It  is  somewhat  soluble  in  water,  one  hundred  liters  of  water 
dissolving  about  three  liters  of  oxygen  under  ordinary  con- 
ditions. The  presence  of  even  a  small  proportion  of  dissolved 
oxygen  is  exceedingly  important.  Fish  die  in  water  con- 
taining no  dissolved  oxygen;  and  the  oxygen  absorbed  by 
water  assists  in  the  decomposition  of  organic  matter  into 
harmless  gases.  Oxygen  is  slightly  heavier  than  air,  its 
density  being  1.105  on  the  air  standard.  This  means  that 
under  the  same  conditions  of  temperature  and  pressure  a 
given  volume  of  oxygen  is  1.105  times  heavier  than  an  equal 
volume  of  air.  One  liter  of  oxygen  weighs  1.429  grams  at  a 
temperature  of  0°  C.  and  under  a  barometric  pressure  of  760 
millimeters.  Subsequently  it  will  be  shown  that  the  volume 


20  INORGANIC   CHEMISTRY 

of  a  gas  depends  upon  the  temperature  and  the  pressure, 
the  normal  or  standard  temperature  and  pressure  being 
respectively  0°  on  the  centigrade  (C.)  thermometer  and  760 
millimeters  (mm.)  on  the  barometer.  Oxygen  like  other 
gases  has  three  physical  states.  Liquid  oxygen  is  pale  blue. 
It  boils  at  —  182.5  °C.  under  the  normal  pressure  of  the  at- 
mosphere; at  this  temperature  its  specific  gravity  is  1.13  (i.e. 
it  weighs  1.13  times  an  equal  volume  of  water).  It  is  mag- 
netic, for  when  a  strong  electromagnet  is  held  near  it,  the 
liquid  is  attracted  by  the  magnet  just  as  iron  filings  are  by 
an  ordinary  magnet.  Solid  oxygen  has  a  whitish  color.  The 
properties  of  oxygen  enumerated  in  this  paragraph  are  physi- 
cal properties,  because  their  exhibition  does  not  involve  a 
transformation  of  the  oxygen  into  a  new  substance. 

The  most  important  chemical  property  of  oxygen  is  the  ease 
with  which  it  combines  or  interacts  with  other  substances. 
It  belongs  to  the  class  of  active  chemical  elements,  and  the 
chemical  changes  it  undergoes  are  attended  by  varied  and 
interesting  physical  changes.  Oxygen  forms  compounds 
with  all  the  other  single  elements  except  fluorine,  bromine, 
argon,  and  helium  (and  the  other  inert  gases  in  the  atmos- 
phere). With  most  of  them  the  union  is  direct ;  i.e.  the  two 
elements  unite  chemically  to  form  a  compound  having  oxy- 
gen and  the  other  element  as  constituents.  This  direct 
combination  is  often  accompanied  by  light  and  heat,  though 
the  temperature  at  which  combination  occurs  varies  be- 
tween wide  limits.  At  the  ordinary  temperature  oxygen 
unites  slowly  with  most  elements,  though  with  phosphorus 
the  chemical  action  is  quite  rapid,  as  may  be  seen  by  the 
glow  and  fumes  when  a  phosphorus-tipped  match  is  rubbed, 
especially  in  the  dark.  Metals,  such  as  lead,  zinc,  and 
copper,  tarnish  or  rust  slowly ;  i.e.  they  combine  slowly  with 
the  oxygen  of  the  air.  The  chemical  activity  of  oxygen  at 
high  temperatures  is  readily  shown  by  putting  burning  or 


OXYGEN  — OZONE  21 

glowing  substances  into  it.  All  burn  vividly  in  oxygen.  It 
interacts  chemically  with  the  substances,  and  the  intensity 
of  the  chemical  energy  is  shown  by  the  heat  and  light  evolved. 
The  vigorous  chemical  activity  of  oxygen  can  be  shown  by 
simple  experiments.  When  a  glowing  stick  of  wood  is  put 
into  oxygen,  the  wood  instantly  bursts  into  flame ;  and  if 
left  in  the  oxygen,  the  wood  continues  to  burn  brightly  until 
the  gas  is  exhausted.  If  faintly  glowing  charcoal  is  put  into 
oxygen,  the  charcoal  glows  vividly.  Sulphur  burns  in  air 
with  a  feeble,  blue  flame,  but  in  oxygen  the  flame  is  much 
larger  and  brighter.  The  flame  in  both  cases  is  accompanied 
by  a  gaseous  product  which  smells  like  a  burning  sulphur 
match.  Iron  wire  does  not  burn  in  air,  but  if  the  end  is 
coated  with  burning  sulphur  and  then  put  into  a  bottle  of 
oxygen,  the  wire  burns  vividly,  throwing  off  a  shower  of 
sparks ;  a  globule  of  white-hot  molten  iron  oxide  is  often  seen 
on  the  end  of  the  wire,  and  sometimes  the  inside  of  the 
bottle  is  coated  with  a  reddish  powder,  which  is  mainly  a 
compound  of  iron  and  oxygen.  Iron  and  oxygen  combine 
at  a  higher  temperature  than  do  sulphur  and  oxygen,  so  the 
heat  from  the  burning  sulphur  is  needed  to  start  the  chemi- 
cal action  between  the  oxygen  and  the  iron.  When  lighted 
magnesium  is  put  into  oxygen,  the  burning  metal  instantly 
becomes  surrounded  with  a  dazzling  flame,  and  burns 
rapidly  to  a  white  powder,  thus  showing  that  the  tempera- 
ture at  which  it  begins  to  combine  with  oxygen  is  much  lower 
than  that  in  the  case  of  iron. 

Test  for  Oxygen.  — The  conspicuous  behavior  of  a  glow- 
ing stick  or  burning  substance  when  put  into  oxygen  enables 
us  to  distinguish  oxygen  from  all  other  gases  except  one 
(see  Nitrous  Oxide).  Thus,  if  a  glowing  stick  is  thrust  suc- 
cessively in  three  bottles  of  gas  and  relights  in  one,  that  gas 
is  probably  oxygen.  This  critical  examination  to  establish 


22  INORGANIC  CHEMISTRY 

the  identity  of  oxygen  is  called  testing  or  making  a  test. 
Each  element  has  properties  which  respond  to  appropriate 
tests.  All  compounds  likewise  behave  in  some  decisive  way 
when  subjected  to  tests.  These  tests  have  been  established 
by  extensive  investigation  and  are  extremely  useful  in  chem- 
istry. 

Chemical  Changes  illustrated  by  the  Chemical  Properties 
of  Oxygen.  —  In  all  the  chemical  changes  described  in  the 
preceding  paragraphs  two  features  are  conspicuous,  viz.  the 
loss  of  physical  properties  of  the  interacting  elements  and 
the  formation  of  new  substances.  The  carbon  (of  the  wood 
and  charcoal),  iron,  sulphur,  and  magnesium  are  no  longer 
recognizable  as  such  in  the  compounds  produced  by  the 
chemical  change.  The  oxygen  likewise  disappears;  for 
although  two  products  are  gases,  their  properties  are  not 
those  of  oxygen.  Thus,  one  gas  (carbon  dioxide)  extinguishes 
flames  and  the  other  (sulphur  dioxide)  has  a  suffocating 
odor.  The  chemical  change  in  the  case  of  the  four  elements 
just  mentioned  is  combination.  The  oxygen  was  added  chem- 
ically to  each  element  and  the  product  was  a  compound  of 
the  two  elements  called  an  oxide.  Their  names  are  carbon 
dioxide,  sulphur  dioxide,  ferrous-ferric  oxide,  and  magnesium 
oxide.  Any  compound  whose  constituents  are  oxygen  and 
one  other  element  is  an  oxide.  (See  Oxides,  below.)  The 
chemical  change  in  the  case  of  wood  is  not  simple  combina- 
tion. Wood  contains  chiefly  a  compound  of  carbon,  hydro- 
gen, and  oxygen,  which  is  first  decomposed  by  heat,  and  then 
the  carbon  and  the  hydrogen  combine  individually  with 
oxygen  as  in  the  other  cases. 

Oxidation.  —  When  sulphur,  iron,  magnesium,  and  carbon 
(in  wood  and  charcoal),  and  other  elements  burn  in  oxygen, 
they  combine  rapidly  with  it.  This  change  is  oxidation, 


OXYGEN  — OZONE  23 

Oxidation  almost  invariably  produces  heat,  though  in  slow 
oxidation  the  heat  is  not  readily  detected  and  is  sometimes 
overlooked.  Very  rapid  oxidation  is  accompanied  by  light. 

The  fact  that  burning  in  its  simplest  aspect  is  a  combining 
with  oxygen  can  be  easily  verified.  .  It  has  been  repeatedly 
shown  that  oxygen  is  one  constituent  of  all  the  products 
formed  by  burning  substances  in  that  gas.  Thus,  carbon 
forms  an  invisible  gas  called  carbon  dioxide,  which  is  a  com- 
pound of  carbon  and  oxygen.  Similarly,  sulphur,  iron,  and 
magnesium  form  compounds  of  these  elements  and  oxygen. 
These  facts  may  be  further  verified  by  a  simple  experiment. 
If  mercury  is  heated,  it  gains  in  weight,  and  red  particles 
collect  on  its  surface;  but  if  it  is  protected  from  the  air  by 
some  coating  and  then  heated,  there  is  no  gain  in  weight 
and  no  evidence  of  the  red  product.  Therefore,  when  the 
exposed  mercury  is  heated,  something  from  the  air  must  be 
added  to  it.  Now,  if  the  red  substance  is  collected  and 
heated  in  a  glass  tube,  mercury  and  oxygen  are  the  only 
products.  Hence,  the  exposed  mercury,  when  heated,  must 
have  combined  with  the  oxygen  of  the  air. 

The  process  of  oxidation  is  not  always  so  simple  as  that 
described  in  the  preceding  paragraph.  Often  one  or  more 
of  the  constituents  of  a  compound  is  seized  upon  by  the 
oxygen,  so  to  speak,  and  converted  into  an  oxide.  Hydro- 
gen is  especially  liable  to  be  thus  removed  from  such  com- 
pounds as  contain  it.  Hence  oxidation,  which  is  primarily 
a  combining  with  oxygen,  is  often  a  complex  process  resulting 
in  the  decomposition  of  a  compound  and  the  formation  of 
one  or  more  oxides. 

Oxidation  is  not  always  rapid  enough  to  produce  light 
and  appreciable  heat.  Thus,  iron  and  other  metals  rust 
and  wood  decays  slowly,  but  both  processes  are  essentially 
oxidation.  Sometimes  slow  oxidation  develops  considerable 
heat.  Thus,  oily  rags,  piles  of  hay,  and  heaps  of  coal  often 


24  INORGANIC  CHEMISTRY 

take  fire  unexpectedly,  because  the  heat  produced  by  the 
continued  oxidation  cannot  escape.  Fires  caused  by  such 
oxidation  are  sometimes  said  to  be  due  to  spontaneous 
combustion. 

Many  substances  give  up  oxygen  readily  and  are  called 
oxidizing  agents.  In  this  class  belong  potassium  chlorate, 
which  is  used  in  fireworks,  and  potassium  nitrate,  which  is 
an  ingredient  of  gunpowder.  In  the  process  of  oxidation, 
oxidizing  agents  lose  oxygen,  and  are  said  to  undergo  reduc- 
tion ;  the  latter  process  will  be  more  fully  described  in  the 
next  chapter. 

Oxides  have  already  been  defined  as  compounds  of  oxygen 
and  one  other  element.  There  are  many  oxides,  and  their 
names  often  express  in  a  general  way  their  composition. 
Oxides  of  different  elements  are  distinguished  by  placing  the 
name  of  the  element  (or  a  slight  modification  of  it)  before  the 
word  oxide,  e.g.  magnesium  oxide,  lead  oxide,  zinc  oxide, 
nitric  oxide,  ferric  oxide.  Sometimes  di-,  or  a  similar 
numerical  syllable,  is  prefixed  to  the  word  oxide,  e.g.  carbon 
monoxide,  manganese  dioxide,  barium  dioxide,  sulphur 
trioxide,  phosphorus  pentoxide.  The  prefix  indicates  the 
proportion  of  oxygen  in  the  oxide. 

Combustion.  —  In  ordinary  language  combustion  means 
fire  or  burning.  Substances  which  kindle  and  burn  readily 
are  called  combustible,  while  those  which  burn  with  diffi- 
culty or  not  at  all  are  called  incombustible.  In  a  limited 
sense,  combustion  is  rapid  oxidation  accompanied  by  heat 
and  often  also  by  light.  Oxygen  is  essential  to  ordinary 
combustion.  If  air  is  excluded  from  a  fire,  the  fire  goes  out. 
When  wood,  coal,  paper,  oil,  or  any  other  combustible  sub- 
stance burns,  the  carbon  (of  which  they  partly  consist)  unites 
with  the  oxygen  of  the  air,  thereby  forming  the  invisible  gas 
carbon  dioxide,  while  the  chemical  change  is  attended  by 


OXYGEN  — OZONE  25 

heat  and  light.  Briefly,  a  burning  substance  is  uniting  rap- 
idly with  oxygen.  But  since  air  is  only  about  one  fifth  oxy- 
gen (the  remainder  being  chiefly  nitrogen,  which  does  not 
support  combustion),  combustion  is  less  rapid  and  hence 
less  vigorous  in  air  than  in  oxygen.  The  compounds  formed 
during  combustion  are  oxides;  in  the  case  of  fuel  they  are 
usually  carbon  dioxide  and  water  —  the  latter  being  hydro- 
gen oxide.  The  temperature  at  which  combustion  takes 
place  varies  between  wide  limits.  Some  substances,  like 
phosphorus  and  gasolene  vapor,  catch  fire  at  a  moderate 
temperature,  while  others  do  not  burn  until  heated  to  the 
very  highest  temperatures.  Each  substance  has  its  own 
kindling  temperature,  i.e.  the  temperature  to  which  it  must 
be  heated  before  it  will  catch  fire,  though  this  temperature 
depends  on  the  form  of  the  substance  as  well  as  on  its  nature 
(see  page  172).  Application  of  this  fact  is  seen  in  the  use 
of  paper  and  kindling  wood  in  starting  a  fire  in  a  stove. 

The  correct  explanation  of  fire,  burning,  and  combustion  was  first 
made  by  Lavoisier  (1743—1794).  For  many  years  chemists  had  be- 
lieved that  all  combustible  substances  contained  a  principle  called 
phlogiston,  and  that  when  a  substance  burned,  phlogiston  escaped. 
Very  combustible  substances  were  thought  to  contain  much  phlogis- 
ton, and  incombustible  substances  no  phlogiston.  This  theory  of 
combustion  was  proposed  by  Becher  (1635-1682)  and  advanced  by 
Stahl  (1660-1734).  Many  famous  chemists  —  Priestley,  Scheele,  and 
Cavendish  —  supported  it.  Lavoisier,  in  1775,  proved  by  his  own 
and  others'  experiments  that  phlogiston  did  not  exist,  and  that 
ordinary  combustion  is  a  process  of  combination  with  "a  certain 
substance  contained  in  the  air."  Soon  after  he  identified  this  sub- 
stance as  oxygen. 

Combustion,  in  a  broad  sense,  is  not  exclusively  oxidation,  but  any 
chemical  change  which  is  attended  by  light  and  heat.  Thus,  hydro- 
gen and  other  elements  burn  in  chlorine,  and  phosphorus  often  pro- 
duces heat  and  light  by  its  combinations  with  other  elements  than 
oxygen.  This  broader  meaning  will  be  illustrated  and  discussed 
later. 


26  INORGANIC  CHEMISTRY 

Relation  of  Oxygen  to  Life.  —  Oxygen  is  essential  to  all 
forms  of  animal  and  plant  life.  If  an  animal  or  a  plant  is 
deprived  of  air,  it  dies.  By  respiration  air  is  drawn  into  the 
lungs,  where  part  of  its  oxygen  is  given  to  the  blood ;  this  oxy- 
gen, which  is  distributed  to  all  parts  of  the  body  by  the  blood, 
oxidizes  the  tissues  of  the  body.  By  this  slow  oxidation 
waste  products  are  formed  and  heat  is  supplied  to  the  body. 
One  of  these  waste  products  is  carbon  dioxide  gas,  which 
with  other  gases  is  exhaled  from  the  lungs.  New  tissue  is 
built  up  from  the  food  we  eat.  The  human  body  resembles 
a  steam  engine.  In  each,  the  oxygen  of  the  air  helps  burn 
fuel  largely  composed  of  carbon.  In  the  engine,  the  products 
escape  through  a  chimney  and  the  heat  produced  is  used  to 
form  steam  which  moves  parts  of  the  machine;  in  the  body, 
the  products  escape  mainly  through  the  lungs  and  the  heat 
keeps  the  body  at  the  temperature  at  which  it  can  best 
perform  its  functions. 

The  vital  necessity  to  fish  of  the  small  proportion  of  oxygen 
dissolved  by  water  should  be  recalled  in  this  connection. 

Oxygen  is  often  administered  to  a  person  who  is  too  weak 
to  inhale  enough  air.  Oxygen  is  also  used  in  various  forms 
of  emergency  breathing  apparatus.  The  pulmotor,  for 
example,  is  a  kind  of  pump  by  which  air  rich  in  oxygen  can 
be  forced  into  the  lungs  at  the  normal  rate  of  respiration ; 
it  is  used  to  resuscitate  persons  who  have  been  overcome  by 
smoke  or  gas  or  who  have  been  rendered  unconscious  by  an 
electric  shock.  An  oxygen-breathing  apparatus  is  used  by 
men  who  must  work  in  places  containing  smoke  or  poisonous 
gases,  e.g.  in  a  mine  after  an  explosion.  The  apparatus  con- 
sists essentially  of  a  helmet  or  a  mouth-breathing  device  con- 
nected by  flexible  tubes  with  a  cylinder  of  compressed  oxygen 
and  a  regenerating  can.  The  apparatus  is  carried  much  like 
a  knapsack.  The  supply  of  oxygen  is  regulated  by  a  valve 
in  the  cylinder,  the  nitrogen  in  the  original  supply  of  air  is 


OXYGEN  —  OZONE  27 

breathed  over  and  over,  and  the  carbon  dioxide  and  water 
vapor  from  the  lungs  are  absorbed  by  potassium  hydroxide 
in  the  regenerating  can.  A  trained  man  equipped  with  this 
apparatus  can  work  as  long  as  two  hours  in  a  vitiated  atmos- 
phere. 

Decay  of  organic  matter  involves  oxidation,  which  is  often 
hastened  by  bacteria.  The  carbon  and  hydrogen  in  the 
organic  matter  form  principally  carbon  dioxide  and  water; 
decay  in  this  respect  resembles  combustion  and  respiration. 

Uses  of  Oxygen.  —  Oxygen  for  commercial  use  is  stored 
under  pressure  in  strong  iron  cylinders.  A  mixture  of  oxy- 
gen and  hydrogen  gas  or  acetylene  gas  if  burned  in  a  suitable 
apparatus  -produces  an  intensely  hot  flame.  The  oxy- 
hydrogen  flame  is  used  to  melt  refractory  metals  and  to  pro- 
duce the  brilliant  light  of  the  stereopticon ;  the  acetylene 
flame  is  used  in  welding  and  also  in  dismantling  iron  and  steel 
structures,  e.g.  bridges,  fences,  and  frames  of  buildings. 

Discovery  of  Oxygen.  —  Oxygen  was  discovered  on  August 
1,  1774,  by  Priestley  (1733-1804).  He  prepared  it  by  focus- 
ing the  sun's  rays  upon  red  mercuric  oxide  by  means  of  "a 
burning  lens  of  twelve  inches  focal  distance."  It  was  inde- 
pendently discovered  by  Scheele  (1742-1786),  a  Swedish 
chemist,  about  the  same  time. 

Priestley  called  the  gas  dephlogisticated  air  because  he 
regarded  it  as  "devoid  of  phlogiston."  Scheele  called  it 
empyreal  air,  i.e.  fire  air  or  fire-supporting  air,  because  it 
assisted  combustion.  Lavoisier,  in  1778,  gave  it  the  name 
oxygen  (from  the  Greek  oxus,  acid,  and  gen,  the  root  of  a  verb 
meaning  to  produce)  because  he  believed  from  his  experiments 
that  oxygen  was  necessary  for  the  production  of  acids  —  a 
view  now  known  to  be  incorrect. 


28  INORGANIC  CHEMISTRY 

PROBLEMS 

1.  Potassium  chlorate  contains  39.18  per  cent  of  oxygen.     How 
many  grams  of  oxygen  can  be  prepared  from  725  gm.  of  potassium 
chlorate  ? 

2.  What  approximate  weight  of  oxygen  can  be  prepared  from  100 
gm.  of  potassium  chlorate  containing  12  per  cent  of  impurity? 

3.  What  is  the  weight  of  (a)  10  1.  of  oxygen,  (6)  75  L,  (c)  500  cc., 
(d)  750  cc.,  (e)  4  1.  (at  standard  temperature  and  pressure)  ? 

4.  A  room  25  m.  long,  17  m.  wide,  and  15  m.  high  is  filled  with 
oxygen.     What  weight  of  gas  does  it  contain  ? 

5.  Perform  the  problems  in  the  Appendix,  §  2. 


OZONE 

Ozone  is  a  gas  related  to  oxygen,  though  its  properties 
differ. 

Formation  and  Preparation.  —  Ozone  is  formed  when 
electric  sparks  pass  through  the  air,  and  is  therefore  pro- 
duced when  electrical  machines  are  in  operation  and  during 
thunder  storms.  Slow  oxidation,  especially  of  moist  phos- 
phorus, produces  ozone.  Its  formation  accompanies  the 
burning  of  hydrogen  (in  oxygen),  and  the  passage  of  elec- 
tricity through  a  solution  of  sulphuric  acid  and  water. 
Ozone  is  prepared  by  subjecting  cold,  dry  oxygen  or  air  to  a 
silent  electric  discharge. 

Properties.  —  Ozone  has  a  peculiar  odor,  suggesting  burn- 
ing sulphur.  The  name  ozone  signifies  smell.  It  has  a 
bluish  color,  and  at  a  low  temperature  condenses  to  a  blue 
liquid.  Liquid  ozone  boils  at  —  119°  C.  under  atmospheric 
pressure.  It  is  a  powerful  oxidizing  agent,  tarnishing  metals, 
bleaching  colored  vegetable  substances,  deodorizing  foul 
animal  matter,  and  corroding  such  substances  as  cork  and 
rubber.  When  heated  to  250°  C.,  or  higher,  it  is  wholly 
changed  into  oxygen.  When  oxygen  is  changed  into  ozone, 


OXYGEN  —  OZONE  29 

it  is  found  that  three  volumes  of  oxygen  yield  two  volumes 
of  ozone ;  and,  conversely,  the  two  volumes  of  ozone,  when 
heated,  become  three  volumes  of  oxygen.  Hence,  volume 
for  volume,  ozone  is  1.5  times  heavier  than  oxygen.  Its 
theoretical  relation  to  oxygen  will  be  subsequently  discussed. 

Uses.  —  Pure  ozone  is  seldom  prepared,  but  air  containing 
ozone  is  used  to  sterilize  drinking  water,  as  a  bleaching  agent, 
and  a  disinfectant. 

EXERCISES  AND  PROBLEMS 

1.  Suggest  experiments  to  show  (a)  that  in  preparing  oxygen 
from  a  mixture  of  potassium  chlorate  and  manganese  dioxide  the 
latter  is  unchanged,  and  (6)  that  air  contains  oxygen. 

2.  Define  and  illustrate  (a)  oxide,  (6)  oxidation,  (c)  oxidizing 
agent. 

3.  Cite  cases  of  spontaneous  combustion. 

4.  Potassium  chlorate  contains  39.18  per  cent  of  oxygen.     How 
many  grams  are  needed  to  prepare  25  1.  of  oxygen  (at  0°  C.  and 
760  mm.)  ? 

5.  Potassium    chlorate   contains   39.18   per   cent   of   oxygen. 
How  many  liters  of  oxygen  (at  0°  C.  and  760  mm.)  can  be  prepared 
from  75  gm.  of  potassium  chlorate  ? 

6.  Mercuric  oxide  contains  7.4  per  cent  of  oxygen.   What  weight 
of  mercuric  oxide  must  be  decomposed  to  yield  25  gm.  of  oxygen  ? 

7.  Mercuric  oxide    contains    7.4    per  cent  of    oxygen.     How 
many  grams  of  mercuric  oxide  will  be  needed  to  prepare  5  1.  of 
oxygen  gas  (at  0°  C.  and  760  mm.)  ? 

8.  Mercuric  oxide  contains  7.4  per  cent  of  oxygen.     How  many 
liters  of  oxygen  gas  (at  0°  C.  and  760  mm.)  can  be  prepared  from 
525  gm.  of  mercuric  oxide  ? 

9.  If  one  third  of  the  oxygen  in  manganese  dioxide  can  be  liber- 
ated by  heating,  what  (a)  weight  and  (6)  volume  of  oxygen  can  be 
obtained  from  25  gm.  of  manganese  dioxide?     (NOTE  :  Manganese 
dioxide  contains  36.78  per  cent  of  oxygen.) 

10.  When  2  gm.  of  a  certain  substance  were  heated,  all  the 
oxygen  which  the  substance  contained  was  given  off,  and  a  residue 
weighing  1.07  gm.  was  left.  Calculate  the  percentage  of  oxygen 
in  the  substance. 


CHAPTER  III 
Hydrogen 

Occurrence. — Free  hydrogen  gas  is  present  in  the  gases 
which  escape  from  volcanoes,  petroleum  wells,  and  natural 
gas  openings.  Artificial  illuminating  gas  contains  consider- 
able hydrogen.  It  is  a  product  of  fermentation  and  decay, 
and  according  to  recent  observations  a  very  small  quantity 
is  present  in  the  atmosphere  of  the  earth.  Enormous  quan- 
tities of  free  hydrogen  exist  in  the  atmosphere  of  the  sun,  and 
during  an  eclipse  of  the  sun  gigantic  streams  of  glowing  hy- 
drogen may  be  seen  shooting  out  from  the  sun's  disk  thou- 
sands of  miles  into  space.  Other  heavenly  bodies  contain 
hydrogen.  Meteorites,  which  come  from  regions  far  beyond 
our  earth,  sometimes  contain  free  hydrogen. 

Combined  hydrogen  is  abundant  and  widely  distributed. 
It  forms  about  one  ninth  (by  weight)  of  water.  Animal  and 
vegetable  substances  contain  hydrogen  in  combination  with 
oxygen  and  carbon,  and  sometimes  nitrogen.  It  is  an  essen- 
tial constituent  of  all  acids.  Combined  with  carbon,  it 
forms  many  gases  and  liquids  called  hydrocarbons,  which 
are  constituents  of  illuminating  gas,  kerosene,  and  gasolene. 
Combined  with  carbon  and  oxygen,  it  forms  such  compounds 
as  sugar,  starch,  paper,  wood,  and  numerous  artificial  prod- 
ucts. With  nitrogen  it  forms  the  familiar  compound,  am- 
monia, and  with  sulphur,  the  bad  smelling  gas,  hydrogen 
sulphide,  which  occurs  in  many  sulphur  springs. 

Preparation.  —  Hydrogen,  like  oxygen,  is  prepared  from 
its  compounds.  This  is  usually  done  by  allowing  certain 

30 


HYDROGEN  31 

metals  and  acids  to  interact.  The  metals  usually  employed 
are  zinc,  iron,  or  magnesium,  and  the  acids  are  dilute  sul- 
phuric acid  or  hydrochloric  acid.  The  hydrogen  comes  from 
the  acid  and  bubbles  through  the  liquid,  when  the  acid  and 
metal  are  mixed.  Hydrogen  is  prepared  in  the  laboratory 
by  putting  the  metal  and  dilute  acid  in  a  glass  vessel  provided 
with  a  delivery  tube  arranged  to  collect  gas  over  water  in  a 
pneumatic  trough.  No  flame  should  be  near  during  the 
performance  of  this  experiment,  because  a  mixture  of  air  and 
hydrogen  explodes  violently  when  ignited.  The  interaction 
of  zinc  and  sulphuric  acid  produces,  besides  hydrogen,  a 
soluble  compound  called  zinc  sulphate. 

Hydrogen  can  also  be  prepared  from  bases  (compounds  of 
hydrogen,  oxygen,  and  a  metal).  Thus,  when  aluminium 
is  boiled  with  sodium  hydroxide,  hydrogen  is  formed. 

Hydrogen  can  be  obtained  from  water  by  allowing  sodium 
or  potassium  to  react  with  it.  If  a  small  piece  of  sodium  is 
dropped  upon  cold  water,  the  sodium  melts  into  a  shining 
globule,  which  spins  about  rapidly  on  the  surface  with  a 
hissing  sound,  and  finally  disappears  after  a  slight  explosion. 
But  when  the  sodium  is  wrapped  in  a  piece  of  tea  lead  pierced 
with  a  few  holes  and  then  dropped  beneath  the  shelf  of  a 
pneumatic  trough  filled  with  water,  the  action  proceeds 
smoothly;  hydrogen  gas  rises  and  displaces  the  water  from 
a  test  tube  or  bottle  supported  over  the  hole  in  the  shelf. 

Hydrogen  can  also  be  prepared  by  the  interaction  of  steam 
—  the  gaseous  form  of  water  —  and  certain  other  metals, 
if  they  are  heated.  This  experiment  was  first  performed  by 
Lavoisier,  in  1783,  while  he  was  studying  the  composition  of 
water.  He  passed  steam  through  a  red-hot  gun  barrel  con- 
taining bits  of  iron,  and  hydrogen  escaped  from  the  tube. 
Since  Lavoisier  was  studying  the  composition  of  water  and 
not  the  properties  of  hydrogen,  he  naturally  thought  of  this 
gas  as  essential  for  forming  water.  So  he  says  in  his  notes, 


32  INORGANIC  CHEMISTRY 

"No  name  appears  to  us  more  suitable  than  that  of  hydro* 
gen,  that  is  to  say,  'generative  principle  of  water.'"  Apart 
from  historical  interest,  this  experiment  has  commercial  im- 
portance, for  if  steam  is  passed  over  white-hot  coal  (instead 
of  iron),  producer  gas  is  formed.  This  mixture  consists  of 
about  one-half  hydrogen,  and  is  used  as  a  source  of  heat  in 
making  steel  and  glass.  If  oil  vapor  is  added  to  this  mixture, 
water  gas  is  formed.  This  is  an  illuminating  gas,  and  is  used 
in  many  cities.  (See  Water  Gas.) 

Hydrogen,  together  with  oxygen,  is  liberated  from  water 
by  passing  a  current  of  electricity  through  water  containing 
a  little  sulphuric  acid.  (See  Chapter  V.) 

Chemical  Changes  illustrated  by  the  Preparation  of  Hy- 
drogen. —  The  preparation  of  hydrogen  by  the  interaction 
of  a  metal  and  an  acid  illustrates  the  third  kind  of  chemical 
change,  viz.  substitution,  or,  as  it  is  sometimes  called,  dis- 
placement or  replacement.  Zinc  and  dilute  sulphuric  acid 
are  usually  used.  The  hydrogen  is  displaced  from  the  acid 
and  the  zinc  takes  its  place ;  i.e.  zinc  is  substituted  chemi- 
cally for  hydrogen.  The  chemical  change  is  not  essentially 
different  from  decomposition  and  combination,  for  we  might 
picture  the  acid  as  decomposing  into  two  fundamental  parts, 
one  part  escaping  as  hydrogen  gas,  the  other  part  (the  sul- 
phur and  oxygen)  at  once  combining  with  the  zinc  to  form 
the  new  compound,  zinc  sulphate.  It  is  better,  however,  to 
regard  substitution  as  a  chemical  change  in  which  one  ele- 
ment replaces  another  in  a  compound,  thereby  producing  a 
different  element  and  a  different  compound.  The  chemical 
change  in  the  preparation  of  hydrogen  from  zinc  and  sul- 
phuric acid  can  be  expressed  by  the  following  equation :  - 

Zinc  +  Sulphuric  Acid  =  Hydrogen  -f-  Zinc  Sulphate. 

The  preparation  of  hydrogen  by  the  interaction  of  water  and 
sodium  is  also  a  case  of  substitution.  Here  the  sodium  dis- 


HYDROGEN  33 

places  one  half  of  the  hydrogen  of  the  water,  thereby  pro- 
ducing free  hydrogen  and  a  new  compound  (sodium  hydrox- 
ide), which  consists  of  sodium,  oxygen,  and  the  rest  of  the 
hydrogen.  The  sodium  hydroxide  dissolves  in  the  water  in 
the  trough.  Briefly,  sodium  is  substituted  for  hydrogen. 
The  change  can  be  expressed  thus :  - 

Sodium  -H  Water  =  Hydrogen  +  Sodium  Hydroxide. 

Physical  Properties.  —  Hydrogen  has  no  taste  or  color. 
The  pure  gas  has  no  odor,  though  hydrogen  as  ordinarily 
prepared  may  have  a  disagreeable  odor,  due  to  impurities. 
Hydrogen  is  the  lightest  known  substance. 
One  liter  of  dry  hydrogen  at  0°  C.  and  760  mm. 
weighs  only  0.08987  gm.  Volume  for  volume, 
hydrogen  is  about  14.4  times  lighter  than  air 
and  about  16  times  lighter  than  oxygen. 

The  extreme  lightness  of  hydrogen  can  be 
easily  shown.  (1)  If  a  wide-mouth  bottle  of 
the  gas  is  left  uncovered  two  or  three  minutes 
and  a  lighted  match  then  dropped  in,  the  match 
will  continue  to  burn.  If  hydrogen  had  been 
present,  it  would  have  combined  with  the 
oxygen  of  the  air  with  a  loud  explosion,  as 
soon  as  the  flame  of  the  match  reached  the 
mixture.  (2)  If  a  bottle  of  hydrogen  is  held  beneath  a  bottle 
of  air  in  the  position  shown  in  Figure  2,  the  gases  exchange 
places,  the  hydrogen,  owing  to  its  lightness,  rising  into  the 
upper  bottle.  Its  presence  there  can  be  readily  shown  by 
inserting  a  lighted  match  into  this  bottle;  if  the  experiment 
has  been  properly  done,  the  hydrogen  will  burn  quickly  and 
quietly,  but  in  most  cases  a  loud  explosion  shows  that  only 
a  part  of  the  hydrogen  has  succeeded  in  entering  the  upper 
bottle.  A  lighted  match  dropped  into  the  other  bottle  re- 
veals air,  or  a  mixture  of  air  and  hydrogen  (if  the  experiment 


34  INORGANIC  CHEMISTRY 

has  been  performed  too  hastily).  (3)  If  a  small  collodion, 
or  rubber,  balloon  is  filled  with  hydrogen  and  then  released, 
it  will  rise  rapidly  into  the  air.  Hydrogen,  because  of  its 
lightness,  is  used  to  fill  balloons;  but  ordinary  illuminating 
gas,  which  is  cheaper,  is  sometimes  used. 

Hydrogen  is  not  very  soluble  in  water,  but  it  is  absorbed 
by  several  metals,  especially  the  rare  metal  palladium.  The 
absorption  of  gases  by  metals  is  called  occlusion.  Only 
about  1.84  1.  of  hydrogen  at  760  mm.  pressure  dissolve  in 
100  1.  of  water  at  20°  C.  Palladium  absorbs  from  300  to  900 
times  its  own  volume  of  hydrogen,  according  to  the  condi- 
tions of  the  experiment.  Platinum,  gold,  and  iron  act  simi- 
larly, though  to  a  less  degree.  Illuminating  gas,  which  con- 
tains considerable  hydrogen,  is  also  absorbed  by  certain 
metals.  Heat  is  developed  by  occlusion.  This  heat  may  be 
sufficient  to  raise  the  metal  to  a  red  heat  and  to  ignite  the  gas 
itself.  One  form  of  self-lighting  gas  burner  acts  on  this 
principle.  Occlusion  is  partly  chemical  and  partly  physical. 

Hydrogen  diffuses  readily;  i.e.  it  quickly  passes  through 
porous  substances,  mixes  with  other  gases  without  stirring 
or  agitating,  and  freely  escapes  into  space  in  all  directions. 
Hydrogen  has  the  highest  rate  of  diffusion,  because  its  den- 
sity is  the  lowest,  it  being  a  general  fact  that  the  rate  of  dif- 
fusion of  a  gas  is  inversely  proportional  to  the  square  root  of 
the  density.  Thus,  the  rate  of  diffusion  of  hydrogen  is 
four  times  that  of  oxygen,  since  the  density  of  oxygen  is  six- 
teen times  that  of  hydrogen.  We  are  largely  indebted  for 
our  knowledge  of  diffusion  to  the  English  chemist,  Thomas 
Graham  (1805-1869). 

Liquid  hydrogen  is  colorless  and  transparent.  It  was  first 
obtained  by  Dewar  in  1898.  At  the  ordinary  pressure  of  the 
atmosphere  liquid  hydrogen  boils  at  —  252.5°  C.  When 
cooled  to  about  —  256°  C.  by  evaporation  under  reduced  pres- 
sure, the  liquid  becomes  a  mass  of  solid  hydrogen;  th$  latter 


SIR    JAMES    DEWAR 


HYDROGEN  35 

is  a  white  froth  if  produced  while  boiling,  and  a  transparent, 
ice-like  solid  if  cooled  when  quiet.  Solid  hydrogen  melts  at 
—  260°  C.  (if  the  pressure  is  58  mm.). 

Hydrogen  is  not  poisonous,  if  pure.  It  does  not  support 
life,  but  a  little  may  be  breathed  without  danger.  When 
the  lungs  are  filled  with  it,  the  voice  becomes  very  thin  and 
shrill. 

Chemical  Properties. — Hydrogen  combines  readily  with  sev- 
eral elements,  especially  chlorine,  oxygen,  sulphur,  and  nitro- 
gen. The  conditions  favorable  for  combination  vary  greatly. 

When  organic  matter  containing  nitrogen  and  sulphur 
decays,  the  products  include  ammonia  and  hydrogen  sul- 
phide. The  former  is  a  compound  of  nitrogen  and  hydrogen, 
and  the  latter  of  sulphur  and  hydrogen.  Hydrogen  and 
chlorine  gases  do  not  unite  in  the  dark,  but  in  the  sunlight 
they  combine  with  explosive  violence  even  at  the  ordinary 
temperature.  Hydrogen  gas  burns  quietly  in  chlorine 
gas.  The  flame  is  bluish  white,  not  very  hot,  and  the 
product  is  hydrochloric  acid  gas  — a  compound  of  hydro- 
gen and  chlorine.  This  burning  of  hydrogen  and  chlorine 
illustrates  the  broader  use  of  the  word  combustion.  No 
oxygen  is  involved.  It  is  a  chemical  change  attended  by 
light  and  heat  and  belongs  to  the  class  of  changes  called 
combination,  since  the  hydrogen  and  chlorine  are  added 
or  combined  chemically. 

Hydrogen  and  oxygen  do  not  unite  at  the  ordinary  tempera- 
ture, but  at  about  750°  C.  the  gases  combine  with  explosive 
violence.  This  temperature  is  provided  by  an  ordinary 
flame  or  red-hot  wire.  A  mixture  of  hydrogen  and  air 
explodes  when  ignited.  Therefore,  the  air  should  be  fully 
expelled  from  the  apparatus  in  which  hydrogen .  is  being 
generated  before  the  gas  is  collected,  and  no  flame,  large 
or  small,  should  be  near.  Neglect  of  these  precautions 


36 


INORGANIC   CHEMISTRY 


has  caused  serious  accidents.  A  small  jet  of  hydrogen, 
however,  burns  quietly  in  air  or  oxygen.  The  flame  is 
almost  invisible  and  very  hot.  Water  is  the  sole  product 
of  the  combustion  of  hydrogen.  The  equation  for  this 
chemical  change  is  :  — 

Hydrogen  +  Oxygen  =  Water. 

These  properties  of  the  hydrogen  flame  can  be  readily 
shown  by  generating  hydrogen  in  a  suitable  apparatus 
and  lighting  the  dry  gas  as  it  issues  from  a  small  opening 
(Fig.  3).  A  platinum  wire  quickly  becomes  red  hot  in  the 


FIG.  3.  —  Apparatus  for  burning  hydrogen.  Acid  is  slowly  introduced 
through  the  funnel  into  the  flask,  which  contains  zinc.  The  liberated  hy- 
drogen is  dried  as  it  passes  through  the  U-tube  containing  calcium  chloride 
and  is  lighted  at  the  platinum  tip  after  all  the  air  has  been  driven  from 
the  apparatus.  The  tip,  which  is  attached  to  the  delivery  tube  by  a  small 
rubber  tube,  is  shown  (about  actual  size)  on  the  left. 

flame,  and  a  dry  inverted  bottle  into  which  the  flame 
is  inserted  becomes  coated  on  the  inside  with  moisture 
condensed  from  the  steam.  The  film  of  water  often  noticed 
on  the  bottom  of  a  vessel  placed  over  a  lighted  gas  range 
or  a  Bunsen  burner  is  formed  by  the  burning  of  the  hydrogen 
and  of  the  hydrogen  compounds  in  the  illuminating  gas. 
The  fact  that  water  is  the  only  product  of  burning  hydro- 
gen was  first  shown  in  1783  by  Cavendish  (1730-1810). 


HYDROGEN  37 

Lavoisier  in  the  same  year  verified  this  fact  and  utilized 
it  to  explain  the  composition  of  water. 

The  temperature  of  the  hydrogen  flame  is  very  high. 
More  heat  is  produced  by  burning  hydrogen  and  oxygen 
than  by  burning  the  same  weight  of  any  other  substance. 

Hydrogen  does  not  support  combustion,  as  the  term  is 
usually  used.  This  property  is  illustrated  by  putting 
a  lighted  taper  into  an  inverted  bottle  of  hydrogen.  The 
taper  ignites  the  hydrogen,  which  burns  at  the  mouth  of  the 
bottle,  but  the  taper  does  not  burn  inside  the  bottle.  Hence, 
hydrogen  burns,  but  does  not  support  combustion.  When 
the  extinguished  taper  is  slowly  withdrawn  through  the 
burning  hydrogen,  however,  it  is  relighted. 

Hydrogen  removes  oxygen  from  compounds  (see  Reduction) . 

Test  for  Hydrogen.  —  The  test  for  hydrogen  is  that  it 
extinguishes  a  flaming  stick  but  is  lighted  at  the  same  time, 
often  with  an  explosion,  and  continues  to  burn  until  the 
gas  is  exhausted. 

The  Oxyhydrogen  Blowpipe  utilizes  the  intense  heat 
produced  by  burning  a  mixture  of  hydrogen  and  oxygen. 
The  apparatus  (Fig.  4) 
consists  of  two  pointed 
metal  tubes.  The  inner 
and  smaller  one  is  for  the 
oxygen,  and  the  outer  Oxyhydrogen  blowpipe  tip. 

and  larger  one  for  the  hydrogen.  Their  pointed  ends  are 
close  together,  and  the  two  gases  mix  as  they  are  forced 
out  of  these  small  openings  by  the  pressure  maintained  in 
the  storage  tanks.  Sometimes  the  tubes  are  separated, 
but  the  gases  flow  from  a  similar  opening.  The  hydrogen 
is  first  turned  on  and  lighted  at  the  pointed  opening,  then 
the  oxygen  is  turned  on  and  the  flow  gradually  regulated 
until  the  flame  is  the  desired  size,  usually  thin,  straight, 


38 


INORGANIC   CHEMISTRY 


and  as  long  as  required.  There  is  no  danger  in  using  the 
blowpipe,  provided  it  does  not  leak  and  the  pressure  of  the 
gases  is  properly  regulated  by  the  stopcocks.  In  the  hot 
flame,  some  metals,  like  silver,  turn  to  vapor;  some,  like 
iron,  burn  brilliantly;  while  others,  like  platinum,  melt. 
When  the  flame  strikes  against  a  piece  of  lime,  the  latter 
becomes  intensely  bright.  Thus  used,  it  is  called  the  lime 
or  calcium  light,  and  is  utilized  in  the  stereopticon.  The  tem- 
perature of  the  oxy hydrogen  flame  is  from  1800  to  2500°  C. 

The  blast  lamp  is  a  modification  of  the  oxyhydrogen 
blowpipe.  The  apparatus  (Fig.  5)  consists  of  two  tubes, 
an  inner  one  for  air  and  an  outer 
one  for  illuminating  gas.  The  air, 
which  is  forced  through  the  apparatus 
by  a  bellows,  provides  oxygen,  while 
the  illuminating  gas  contains  hydro- 
gen and  other  combustible  gases. 
The  mixture  burns  at  the  opening 
of  the  tubes  with  a  colorless  or  bluish 
flame,  which  is  hotter  than  the  Bun- 
sen  flame  (the  usual  source  of  heat 
for  chemical  experiments) .  The  shape 
and  size  of  the  flame  are  easily  regu- 
lated by  stopcocks.  The  blast  lamp  is  used  as  a  source  of 
heat  for  many  operations  in  the  laboratory,  especially  in 
chemical  analysis. 

Reduction.  —  Hydrogen  not  only  combines  energetically 
with  free  oxygen,  but  it  also  withdraws  oxygen  from  com- 
pounds. The  chemical  removal  of  oxygen  is  called  re- 
duction. Hydrogen  is  a  vigorous  reducing  agent,  just  as 
oxygen  is  an  energetic  oxidizing  agent.  When  oxides  of 
certain  metals  are  heated  in  a  current  of  hydrogen,  the  oxy- 
gen of  the  oxide  is  removed  and  combines  with  the  hydrogen 


FIG.  5.  —  Blast  lamp. 


HYDROGEN  39 

to  form  water;  the  metal  is  left  uncombined.  Thus,  by 
heating  lead  oxide  in  hydrogen,  water  and  metallic  lead  are 
produced.  Chemically  speaking,  the  lead  oxide  is  reduced 
by  the  hydrogen.  The  chemical  change  is  substitution 
(the  hydrogen  being  substituted  chemically  for  the  metal), 
and  it  can  be  expressed  thus  :  — 

Hydrogen  +  Lead  Oxide  =  Water  +  Lead. 

This  chemical  change  can  also  be  interpreted  from  the 
standpoint  of  oxidation,  because  the  hydrogen  is  oxidized 
to  water  at  the  same  time  the  lead  oxide  is  reduced.  In 
fact,  the  processes  of  reduction  and  oxidation  are  closely 
related  and  usually  occur  in  the  same  chemical  change; 
either  one  may  be  emphasized  in  interpreting  the  change. 
It  is  preferable,  however,  at  this  stage  to  define  reduction 
as  the  removal  of  oxygen  from  a  compound,  posl^mg  the 
details  of  the  process  until  more  facts  are  availaW^l^n  its 
simplest  form,  reduction  is  the  opposite  of  oxidatic 


Discovery  of  Hydrogen.  —  Paracelsus  in  the  sixteenth 
century  obtained  hydrogen  by  the  interaction  of  acids  and 
metals.  It  was  identified  as  an  element  in  1766  by  Caven- 
dish, who  called  it  inflammable  air.  The  name  hydrogen, 
given  to  it  by  Lavoisier,  in  1783,  is  derived  from  the  Greek 
word  hudor,  water,  and  gen,  the  root  of  a  verb  meaning  to 
produce.  (See  Preparation  of  Hydrogen,  third  paragraph.) 

PROBLEMS 

1.  What  volume  does  5  gm.  of  hydrogen  occupy  (at  0°  C.  and 
760  mm.)  ? 

2.  The  density  of  chlorine  is  nearly  thirty-six  times  that  of  hydro- 
gen.    Compare  its  rate  of  diffusion  with  that  of  hydrogen. 

3.  A   certain  gas  passes  through  a  porous  partition  2.5   times 
slower  than  hydrogen.     What  is  its  density  ? 

4.  How  many  times  heavier  than  a  liter  of  hydrogen  is  a  liter  of 
oxygen,  both  being  dry  and  under  standard  conditions'? 


40  INORGANIC  CHEMISTRY 

5.  What  is  the  weight  of  (a)  8000  cc.  of  dry  hydrogen  gas  at  0° 
C.  and  760  mm.  ?     (6)  Of  1800  cc.  ?     (c)  Of  9  1.  ? 

6.  The  standard  pressure  at  which  a  gas  is  measured  is  760 
mm.     Express  the  same  in  inches. 

7.  If  sulphuric  acid  contains  2.04  per  cent  of  hydrogen,  how 
many  liters  of  hydrogen  (at  0°  C.  and  760  mm.)  can  be  obtained 
from  137  gm.  of  sulphuric  acid  ? 

8.  Water  contains  11.18  per  cent  of  hydrogen.     How  many 
(a)  cc.  and  (6)  gm.  of  the  gas  can  be  prepared  from  1  1.  of  water? 

9.  What   is  the  weight  in  gm.  of  (a)  50  1.  of  hydrogen  gas? 
(6)  50,000  cc.  ?     (c)  50  cdm.  ? 

10.  (a)  72  1.  of  hydrogen  gas  (at  0°  C.  and  760  mm.)  weigh 
how  many  grams,   (6)  centigrams,   (c)  milligrams,   (d)  kilograms, 
(e)  decigrams? 

11.  How  many  grams  does  a  cubic  meter  of  hydrogen  weigh? 

12.  (a)  How  many  cc.  (at  0°  C.   and  760  mm.)  will  75  gm.  of 
hydrogen  occupy  ?     (6)  How  many  gm.  will  75  cc.  weigh  ? 

13.  A  cylindrical  tank  1.5  m.  long  and  30  cm.  in  diameter  is 
filled  with  hydrogen   (at  0°  C.  and  760  mm.),     (a)  What  is  the 
weight  of  the  gas  in  gm.  ?     (6)  How  many  gm.  of  water  is  needed 
to  prepare  the  gas?      (c)  How  many  kg.  of  sulphuric  acid?     (d) 
How  many  dg.  of  sodium  hydroxide  (containing  1.75  per  cent  of 
hydrogen) ? 


CHAPTER  IV 
Some  Properties  of  Gases 

Introduction.  —  Several  elements  and  many  compounds 
are  gases  or  can  be  readily  changed  into  gases.  Elementary 
gases,  besides  oxygen  and  hydrogen,  are  chlorine  and  nitrogen; 
compound  gases  are  carbon  dioxide,  hydrochloric  acid, 
ammonia,  and  sulphur  dioxide.  Air  is  a  mixture  of  several 
gases,  but  it  behaves  like  a  single  gas.  Water  is  readily  con- 
verted into  the  gaseous  state,  which  is  familiar  as  steam  and 
water  vapor. 

The  properties  of  all  gases  vary  with  the  temperature 
and  the  pressure  to  which  they  are  subjected.  Thus, 
oxygen  gas  becomes  liquid  oxygen  at  a  low  temperature, 
while  liquid  water  is  constantly  changing  into  water  vapor. 
The  most  common  and  conspicuous  change,  however,  is  an 
alteration  in  volume  whenever  pressure  and  temperature 
are  varied. 

Relation  of  Gas  Volumes  to  Temperature  and  Pressure.  — 

The  actual  volume  occupied  by  a  gas  depends  upon  the 
temperature  and  pressure  prevailing  at  the  time  of  obser- 
vation. The  volume  expands  with  rise  of  temperature  or 
with  decrease  of  pressure;  it  contracts  with  fall  of  tempera- 
ture or  with  increase  of  pressure.  In  general,  if  we  cool  a 
gas  or  subject  it  to  a  greater  pressure,  it  shrinks  ;  if  we  heat  a 
gas  or  subject  it  to  a  lower  pressure,  it  expands.  By  common 
consent,  the  normal  or  standard  temperature  is  zero  degrees 
on  the  centigrade  thermometer  (or  briefly  0°  C.),  and  the 
normal  or  standard  pressure  is  the  pressure  indicated  by  the 

41 


42  INORGANIC  CHEMISTRY 

barometer  when  the  mercury  column  is  760  millimeters  high 
(or  briefly  760  mm.).  Under  these  conditions,  which  are 
called  standard  conditions,  a  liter  of  dry  oxygen  gas  weighs 
1.429  gm.  But  at  another  temperature  or  pressure  the 
liter  would  contain  a  different  quantity  of  oxygen  gas,  and 
would  therefore  have  a  different  weight.  For  example,  if 
the  pressure  is  increased,  the  volume  becomes  less,  more  gas 
must  be  added  to  bring  the  volume  up  to  a  liter,  and  this 
second  liter  of  oxygen  would  weigh  more  than  1.429  gm. 
That  is,  a  liter  vessel,  when  full,  always  contains  a  liter, 
but  the  weight  of  the  contents  varies  with  the  quantity  of 
gas  contained  in  this  volume.  Clearly,  if  we  wish  to  com- 
pare the  weights  of  gases  by  means  of  their  volumes,  we  must 
know  the  conditions  under  which  the  volume  is  measured. 
As  we  shall  subsequently  see,  the  comparison  of  weights  of 
gases  is  a  frequent  and  highly  important  operation  in  chemis- 
try. Some  method  is  necessary,  therefore,  to  permit  this 
comparison.  If  all  gases  could  be  measured  at  0°  C.  and 
760  mm.,  their  volumes  would  be  comparable  and  the 
weights  deduced  or  obtained  directly  from  these  volumes 
would  be  a  true  measure  of  the  actual  quantity  of  the -gases 
in  the  observed  volumes.  But  it  is  experimentally  incon- 
venient to  measure  gases  at  0°  C.  and  760  mm.  So  it  is 
customary  to  measure  the  volume  under  the  conditions 
existing  at  the  time  of  the  experiment,  and  then  reduce  the 
observed  volume  to  the  volume  it  would  occupy  under  stand- 
ard conditions.  This  mathematical  reduction  is  performed 
by  applying  two  laws,  —  the  law  of  Charles  and  the  law  of 
Boyle. 

Relation  of  the  Volume  of  a  Gas  to  Changes  in  Tempera- 
ture. —  It  has  been  found  by  experiment  that  all  gases 
under  constant  pressure  expand  or  contract  equally  for 
equal  changes  of  temperature.  If  the  volume  of  a  gas  is 


SOME  PROPERTIES  OF  GASES 


43 


measured  at  0°  C.,  the  gas  expands  or  contracts  1/273  of  this 
volume  for  a  rise  or  fall  of  one  degree.  That  is,  273  vol- 
umes at  0°  C.  become  274  at  1°  C.,  275  at  2°,  280  at  7°,  272 
at  -1°,  270  at  -3°,  or  in  general,  (273  +  0  volumes  at 
t°  (i.e.  at  any  temperature).  The  statement  above  summa- 
rizing the  relation  between  gas  volumes  and  temperature 
is  known  as  the  Law  of  Charles.  It  applies  accurately 


125 

- 

400 

100 

.- 

375 
373 

75 

.  - 

350 

50 

.- 

325 

25 

.- 

300 

0 

.; 

275 
273 

!-25 

.- 

250  g 

! 

>-50 

.- 

225  v 

0 

-75 

200  £ 

-100 

.- 

175  | 

'-125 

.- 

150 

-150 

125 

-175 

.- 

100 

-200 

75 

-225 

.- 

50 

-250 

25 

-273 

0 

CENTIGRADE 

200 

100 

50 

25 
0 

-25 

-50 

-100 

-273 


ABSOLUTE 
473 
373 
323 
298 
273 
248 
223 
173 
0 


FIG.  6.  —  Centigrade  and  absolute  thermometer  scales  (left)   and   some 
equivalent  degrees  (right). 

only  to  the  temperature  which  would  ordinarily  be  used,  not 
to  extreme  temperatures  nor  to  temperatures  near  the  point 
at  which  the  gas  liquefies.  Thus,  according  to  the  law,  if 
a  gas  could  be  cooled  to  —  273°  C.,  its  volume  would  become 
zero !  But  this  low  temperature  has  never  been  reached, 
and  even  if  it  could  be,  all  gases  (except  possibly  helium) 


44  INORGANIC  CHEMISTRY 

become  solid  before  reaching  this  temperature.  Similarly, 
many  gases  dissociate  or  decompose  at  very  high  tempera- 
tures, and  all  deviate  from  the  law  when  about  to  change 
from  the  gaseous  to  the  liquid  state.  This  point  (—  273°  C.) 
on  the  thermometer  is  called  absolute  zero,  and  a  scale 
starting  at  273°  below  0°  C.  is  called  the  absolute  thermom- 
eter scale.  The  relation  between  the  two  scales  is  shown 
in  Figure  6. 

An  examination  of  these  scales  shows  that  absolute 
degrees  are  numerically  greater  by  273  than  the  correspond- 
ing centigrade  degrees.  This  relation  is  sometimes  expressed 
by  the  statement:  - 

To  convert  centigrade  into  absolute  add  273. 

The  application  of  the  law  of  Charles  to  the  reduction 
of  a  gas  volume  to  the  volume  it  would  occupy  at  0°  C.  can 
be  readily  understood  by  an  example.  Suppose  10  cc. 
of  dry  oxygen  gas  at  15°  C.  are  to  be  reduced  to  the  volume 
occupied  at  0°  C.  The  corrected  volume,  as  it  is  often  called, 
can  be  found  by  two  processes.  (1)  The  first  method  utilizes 
the  ordinary  centigrade  scale.  Let  the  volume  at  0°  C.  be 
represented  by  273.  Now  since  the  volume  of  a  gas  at  0°  C. 
expands  1/273  for  each  degree  through  which  it  is  heated, 
the  volume  of  oxygen  at  15°  C.  would  be  represented  by 
273  +  15.  But  273  +  15  and  273  are  in  the  same  ratio  as  10 
(the  known  volume  at  15°  C.)  and  X  (the  unknown  volume 
at  0°  C.).  Therefore  we  can  state  these  relations  in  a 
proportion,  thus :  — 

273  +  15:  273::10:X;  X  =  9.479  cc. 

Therefore,  10  cc.  of  oxygen  at  15°  C.  occupy  9.479  cc.  at 
0°  C.  Since  t  can  be  substituted  for  any  temperature 
(above  or  below  0°  C.),  the  general  form  of  the  proportion 
can  be  written :  — 

273  + 1 :  273  : :  known  vol.  :  vol.  at  0°  C. 


SOME   PROPERTIES   OF  GASES 


45 


(2)  The  second  method  uses  the  absolute  scale  and  is 
based  on  the  fact  that  the  volumes  of  a  gas  at  different 
temperatures  vary  as  their  absolute  temperatures.  Suppose 
we  have  273  cc.  of  a  gas  at  0°  C.  Since  it  expands  1/273 
of  its  volume  at  0°  C.  for  each  degree  of  increase  in  tem- 
perature, and  contracts  1/273  for  each  degree  of  decrease, 
its  volumes  would  be  as  follows  for  certain  centigrade 
temperatures :  — 


VOLUME  IN  cc. 


CENTIGRADE  TEMPERATURES 


373 
323 
273 
223 
173 


100 

60 

0 

-50 

-100 


Comparing  these  values  with  the   corresponding  absolute 
temperatures,  we  have  the  following  relations :  — 


VOLUME  IN  cc. 

^ 

CENTIGRADE  TEMPERATURE 

ABSOLUTE  TEMPERATURE 

373 

100 

373 

323 

50 

323 

273 

0 

273 

223 

-50 

223 

173 

-  100 

173 

It  is  clear  from  the  first  and  third  columns  that  the  volumes 
and  absolute  temperatures  are  numerically  the  same.  Hence 
in  reducing  gas  volumes  to  those  occupied  at  0°  C.  by  the 
absolute  method,  there  are  three  steps,  (a)  Convert  the 
observed  centigrade  temperature  into  absolute  temperature 
by  adding  273.  (6)  Make  273  the  numerator  of  a  fraction 
and  the  sum  found  in  (a)  its  denominator.  (c)  Multiply 


46  INORGANIC  CHEMISTRY 

the  given  or  observed  volume  by  the  fraction  formed  in  (6). 
Using  the  problem  given  above,  the  absolute  process  becomes 
(a)  273  +  15  =  288,  (6)  273/288,  (c)  10  x  273/288  =  9.479. 
The  solution  of  a  few  problems  by  either  of  these  methods 
will  fix  in  mind  the  important  relation  of  gas  volumes  to 
changes  in  temperature. 

Relation  of  Gas  Volumes  to  Changes  in  Pressure.  —  It 
has  been  found,  by  experiment,  that  the  volume  of  a  gas  at  a 
constant  temperature  is  inversely  proportional  to  the  pres- 
sure. This  statement  is  the  Law  of  Boyle,  and  was  an- 
nounced by  him  in  1660.  This  law,  like  the  law  of  Charles, 
applies  accurately  only  to  the  pressures  ordinarily  used,  but 
the  general  numerical  relation  between  volume  and  pressure 
as  stated  above  is  sufficiently  accurate  for  most  purposes. 
Boyle's  law  means  that  the  greater  the  pressure,  the  less  the 
volume,  and  vice  versa.  The  normal  or  standard  pressure 
is  the  mean  atmospheric  pressure;  this  is  equal  to  the 
pressure  of  a  column  of  mercury  760  millimeters  high. 
Briefly,  the  normal  pressure  is  760  mm.  Most  gases  under 
examination  are  confined  over  water  or  some  other  liquid 
whose  surface  is  exposed  to  the  atmosphere,  and  since 
atmospheric  pressure  is  transmitted  through  the  liquid  to 
the  gas,  the  pressure  which  the  gas  is  under  is  found  by 
reading  the  pressure  recorded  by  the  barometer  at  the  time 
the  gas  volume  is  read.  The  reduction  of  the  observed 
volume  to  the  volume  it  would  occupy  at  760  mm.  is  per- 
formed by  applying  Boyle's  law.  An  actual  case  will  make 
the  process  clear.  Suppose  we  have  25  cc.  of  dry  oxygen 
gas  at  775  mm.  and  wish  to  know  its  volume  at  760  mm.  Two 
processes  —  not  essentially  different  —  can  be  used.  (1)  Ac- 
cording to  Boyle's  law,  gas  volumes  are  inversely  pro- 
portional to  the  pressures;  i.e.  the  observed  pressure  bears 
the  same  relation  to  the  normal  pressure  as  the  normal 


SOME   PROPERTIES   OF  GASES  47 

volume  bears  to  the  observed  volume.  This  general  relation 
applied  to  the  problem  becomes  — 

775:760::X:25;  X  =  25.49. 

Therefore,  25  cc.  of  oxygen  at  775  mm.  occupy  25.49  cc.  at 
760  mm.  (2)  The  other  method  involves  two  steps,  (a)  Make 
760  the  denominator  of  a  fraction  and  the  observed  pressure 
its  numerator.  (6)  Multiply  the  observed  volume  by  the 
fraction  formed  in  (a).  Using  the  same  problem,  the  second 
process  becomes  — 

(a)   775/760,     (6)   25x775/760  =  25.49. 

Behavior  of  Gas  Volumes  under  Simultaneous  Action  of 
Heat  and  Pressure.  —  Heat  and  pressure  act  independently 
upon  a  gas.  That  is,  it  is  immaterial  whether  a  gas  is  sub- 
jected to  heat  and  to  pressure  at  the  same  time  or  in  succes- 
sion. The  final  volumes  are  equal  in  either  case.  Since  both 
heat  and  pressure  are  factors  in  the  cause  of  the  changing 
gas  volume,  it  is  convenient  to  reduce  a  gas  volume  to  the 
standard  volume,  so  to  speak,  by  a  single  operation.  Thus, 
if  the  observed  volume  is  25  cc.,  the  temperature  15°  C., 
and  the  pressure  775  mm.,  the  standard  volume  is  found 
thus  :  — 

25  X  775/760  X  273/288  =  X. 

Sometimes  the  reduction  to  standard  conditions  is  per- 
formed by  substituting  the  observed  values  in  the  formula: 


760  [1  +  (.00366  X  0] 

In  the  formula  which  merely  involves  the  mathematical 
operations  just  described,  V  =  final  volume,  V1  =  observed 
volume,  P'  =  observed  pressure,  t  =  observed  temperature, 
and  .00366  =  1/273.1 

1  The  method  of  deducing  this  formula  is  given  in  the  author's  "  Experi- 
mental Chemistry,"  pp.  361-363. 


48  INORGANIC  CHEMISTRY 

It  should  be  noted  that  the  reduction  of  gas  volumes  to 
standard  conditions  is  a  mathematical  process,  and  does 
not  imply  that  the  gas  itself  must  actually  be  subjected  to 
these  conditions.  (See  Laboratory  Manual,  Appendix  B.) 

Weight  of  a  Liter  of  Oxygen  Gas.  — As  already  stated,  the 
weight  of  a  liter  of  oxygen  is  1.429  gm.  at  0°  C.  and  760  mm. 
Its  weight  would  be  different  at  any  other  temperature  and 
pressure  (unless  the  effect  of  heat  and  pressure  balanced 
each  other).  This  value  (1.429)  is  found  by  an  experiment 
involving  several  steps,  (a)  Oxygen  is  generated  from  a 
mixture  of  potassium  chlorate  and  manganese  dioxide,  and 
the  quantity  liberated  is  found  by  subtracting  the  weight 
of  the  oxygen  generator  after  the  experiment  from  its  original 
weight;  suppose  the  weight  of  liberated  oxygen  is  2.312  gm. 
(6)  The  oxygen  is  collected,  its  volume  noted,  and  the  tem- 
perature and  pressure  also  read;  suppose  the  volume  of 
oxygen  is  1.75  1.,  the  temperature  is  19°  C.,  and  the  pressure 
is  755  mm.  (c)  The  observed  volume  is  reduced  to  the 
volume  at  standard  conditions,  thus :  — 

1.75  x  755/760  x  273/292  =  1.62. 

(d)  The  weight  of  one  liter  at  0°  C.  and  760  mm.  is  then 
found  to  be  1.427  gm.  by  dividing  2.312  (the  weight  of  the 
oxygen)  by  1 .62  (the  corrected  volume  of  the  oxygen),  thus :  — 

2.312 -j- 1.62  =  1.427. 

Very  accurate  experimental  work  involving  precautions  not 
mentioned  above  yields  the  value  1.429  gm.  as  the  exact 
weight  of  one  liter  of  oxygen. 

Densities  of  Gases.  —  The  weight  of  a  liter  of  many  gases 
can  be  found  by  a  method  similar  to  that  used  for  oxygen. 
The  values  obtained  are  comparable  because  the  volumes 
are  corrected  for  temperature  and  pressure  (as  reduction 


SOME   PROPERTIES   OF  GASES  49 

to  standard  conditions  is  sometimes  designated).  Com- 
parison of  the  weight  of  a  liter  of  different  gases  reveals 
interesting  relations.  Thus,  a  liter  of  air  under  standard 
conditions  weighs  1.293  gm.  Now  if  we  divide  1.429  by 
1.293,  the  quotient  (1.105)  shows  that  oxygen  is  1.105  times 
heavier  than  air.  If  we  divide  1.429  by  .08987,  which  is 
the  weight  of  a  liter  of  hydrogen  at  0°  C.  and  760  mm., 
the  quotient  (15.9)  shows  that  oxygen  is  15.9  times  heavier 
than  hydrogen.  This  number  (15.9)  is  the  density  of 
oxygen  on  the  hydrogen  standard.  Density  is  the  relative 
weight  of  equal  and  comparable  volumes.  Sometimes 
density  is  defined  as  the  weight  of  a  body  in  grams  divided 
by  its  volume  in  cubic  centimeters.  Thus,  according  to 
the  latter  definition  the  density  of  oxygen  is  1.429  -5-  1000  = 
.001429.  But  since  in  actual  practice  we  more  often  com- 
pare gas  volumes  with  each  other,  especially  with  hydrogen 
or  air,  the  former  definition  (i.e.  relative  weights  of  equal 
volumes)  is  more  convenient.  Important  deductions  are 
made  from  the  densities  of  gases,  as  will  appear  in  a  sub- 
sequent chapter. 

PROBLEMS 

1.  Reduce   the   following  volumes   to   the   volume   occupied   at 
0°  C. :  (a)  173  cc.  at  120°  C.,  (6)  466  cc.  at  14°  C.,  (c)  706  cc.  at  15°  C., 
(d)  25  cc.  at  27°  C. 

2.  Reduce   the   following   volumes   to   the   volume   occupied   at 
760  mm.:   (a)  200  cc.  at  740  mm.,  (6)  25  cc.  at  780  mm.,  (c)  467  cc. 
at  756  mm.  Ans.     (a)  194.7,  (6)  25.65,  (c)  464.54. 

3.  Reduce  the  following  to  the  volume  at  standard  conditions: 
(a)    147  cc.  at  570  mm.  and   136.5°  C.,  (6)  320  cc.  at  950  mm.  and 
91°  C.,  (c)  480  cc.  at  380  mm.  and  68.25°  C.,  (d)  25  cc.  at  780  mm. 
and  27°  C.,  (e)  14  cc.  at  763  mm.  and  11°  C. 

4.  (a)  The  temperature  of  a  gas  is  18°  C.     At  what  temperature 
would  its  volume  be  doubled?     (6)  A  gas  measures  195  cc.  at  740 
mm.     What  is  the  pressure  when  the  volume  is  295  cc.  ? 


CHAPTER  V 


General  Properties  of  Water 

WATER  is  a  compound  of  hydrogen  and  oxygen,  and  is 
worthy  of  extensive  study  because  of  its  indispensable  re- 
lation to  life,  characteristic  properties,  and  numberless  uses. 

Occurrence  in  Nature.  —  Water  in  the  form  of  vapor  is 
always  present  in  the  atmosphere.  Evaporation  is  con- 
stantly taking  place  from  the  surface  of  the  ocean  and  other 
bodies  of  water,  from  the  moist  earth,  from  the  bodies  of 
animals,  and  from  plants.  This  vapor  condenses,  and  appears 
as  clouds,  mist,  fog,  snow,  hail,  dew,  and  frost. 

In  the  liquid  state  water  occurs  in  vast  quantities.  About 
three  fourths  of  the  surface  of  the  globe  is  water.  Soil  and 
porous  rocks  hold  considerable,  and  plants  and  animals  con- 
tain a  large  proportion.  Certain  substances,  which  are 
apparently  dry,  really  retain  much  water.  Thus,  in  a  ton 
of  clover  hay  there  are  upwards  of  200  Ib.  of  water.  Many 
common  foods  consist  largely  of  water,  as  may  be  seen  by 
the  following  — 

TABLE  OF  THE  PROPORTION  OF  WATER  IN  FOOD 


FOOD 

PER  CENT 
OK  WATER 

FOOD 

PER  CENT 
OF  WATER 

Cod  

82.6 

Tomatoes     .    .    , 

94  3 

Beef  

61.9 

Apples     

84.6 

Lobster      

79.2 

Strawberries     .    .    ,    . 

904 

73.7 

Watermelon     .... 

92  4 

94.0 

Milk    

87  0 

Potatoes    

78.3 

Cheese     

28  to  72 

Cucumbers    

95.4 

White  Bread    .... 

35.3 

50 


GENERAL  PROPERTIES  OF  WATER      51 

The  human  body  is  nearly  70  per  cent  water,  and  during 
a  year  the  average  man  drinks  about  half  a  ton. 

Water  in  the  form  of  ice  permanently  covers  the  coldest 
parts  of  the  surface  of  the  earth,  e.g.  the  polar  regions  and 
the  summits  of  high  mountains. 

Functions  of  Water  in  Nature.  —  Since  water  is  the  only 
liquid  occurring  in  large  quantities  on  the  earth's  surface, 
it  is  the  most  effective  agent  of  erosion.  It  cuts  away  the 
earth's  crust,  and  transports  the  material  from  higher  to 
lower  levels,  or  washes  it  ultimately  into  the  ocean.  Acting 
in  conjunction  with  carbon  dioxide  gas,  it  decomposes  the 
rocks,  changing  them  into  clay,  sand,  and  substances  which 
make  the  soil  productive.  The  cycle  of  changes  from  liquid 
to  vapor  and  vapor  to  liquid  exerts  a  marked  influence  on 
the  distribution  of  heat  and  moisture  upon  the  earth's 
surface,  i.e.  on  climate. 

Water  dissolves  many  solids  and  gases,  and  is  constantly 
extracting  from  the  rocks  and  soil  their  soluble  constituents, 
some  of  which  serve  for  the  nutrition  of  plants,  though 
the  larger  part  passes  on  to  the  ocean.  The  latter  thus 
becomes  a  vast  reservoir  of  water  containing  salt  and  other 
mineral  matter  obtained  from  the  earth's  crust.  In  the 
vital  processes  of  animals  and  plants  water  helps  change 
the  food  into  a  condition  fit  for  distribution  and  assimilation. 

Industrial  Applications.  —  Besides  the  universal  use  of 
water  as  a  beverage,  it  is  applied  to  an  endless  variety  of 
useful  and  convenient  purposes.  It  has  always  been  man's 
beast  of  burden.  It  is  the  vehicle  for  transferring  me- 
chanical energy  to  water  wheels  —  an  application  now  being 
made  on  a  vast  scale  for  generating  electricity.  It  utilizes 
by  its  peculiar  properties  the  energy  in  fuel  by  means  of 
the  steam  engine.  It  is  the  highway  for  transportation  on 
the  largest  scale  by  ocean,  rivers,  lake,  and  canal.  It  is 


52  INORGANIC   CHEMISTRY 

the  vehicle  for  the  distribution  of  heat  by  hot  water  and 
steam.  It  is  the  indispensable  solvent  in  metallurgy,  in 
the  manufacture  of  many  chemicals,  and  in  such  industries 
as  soap  making,  bleaching,  brewing,  dyeing,  and  tanning; 
it  is  necessary  wherever  mortar  and  cement  are  used.  Man's 
work  would  be  stopped  in  a  thousand  other  ways  were  he 
deprived  of  water. 

Natural  Waters.  —  Water  is  never  found  pure  in  nature. 
Even  rain  water,  which  is  the  purest  natural  water,  con- 
tains gases  and  dust  washed  from  the  air.  When  rain 
strikes  the  ground,  it  begins  at  once  to  take  up  impurities 
from  the  rocks  and  soil.  Some  of  the  water  flows  along 
the  surface,  becoming  more  and  more  impure,  and  finally 
reaches  the  ocean.  But  25  to  40  per  cent  of  the  annual 
rainfall  in  temperate  regions  soaks  into  the  ground  and 
percolates  through  the  soil  at  an  estimated  rate  of  .2  to 
20  feet  a  day.  On  its  journey  underground  the  water  loses 
most,  often  all,  of  its  organic  matter  (i.e.  vegetable  or  animal 
matter  or  products  of  their  decomposition),  but  it  dissolves 
mineral  matter  and  gases.  The  mineral  matter  is  usually 
common  salt  and  compounds  of  calcium  and  magnesium ; 
the  most  common  gas  is  carbon  dioxide.  If  the  amount 
of  dissolved  matter  in  spring  water  is  large,  or  the  kind  of 
matter  is  so  unusual  as  to  give  the  water  a  marked  taste  or 
medicinal  properties,  the  water  is  called  mineral  water.. 
Water  containing  calcium  and  magnesium  compounds  is 
hard,  but  in  soft  water,  such  as  rain  water,  these  compounds; 
are  absent. 

There  are  several  hundred  mineral  springs  in  the  United  States. 
Those  having  a  high  temperature  are  called  thermal,  as  at  Hot  Springs, 
Arkansas.  Many  contain  a  large  proportion  of  common  -salt,  as  at 
Saratoga,  New  York.  Others  contain  alkaline  substances  and  carbon 
dioxide  gas,  e.g.  Vichy  and  Apollinaris  water.  Sulphur  springs  con- 
tain solid  or  gaseous  compounds  of  sulphur, — or  both, — and  have 


GENERAL  PROPERTIES  OF  WATER 


53 


valuable  medicinal  properties.  Some,  like  Hunyadi,  are  bitter;  but 
others,  especially  those  in  New  York  State,  which  contain  gaseous 
sulphur  compounds,  have  a  sweet  taste  but  an  unpleasant  odor.  Cha- 
lybeate waters  contain  soluble  iron  compounds.  Many  mineral 
waters  contain  calcium  and  magnesium  compounds,  and  a  few  con- 
tain alum  and  lithium  compounds.  Most  natural  mineral  waters 
contain  traces  of  a  large  number  of  different  substances.  Many  com- 
mercial mineral  waters  have  indifferent  medicinal  value. 

River  water  obviously  contains  the  impurities  brought 
by  springs  and  the  surface  water;  it  is  also  often  made  very 
impure  by  decaying  animal  and  vegetable  matter,  which 
has  been  purposely  or  accidentally  introduced,  especially 
if  the  river  passes  through  a  thickly  settled  region.  A 
sluggish  river  is  more  apt  to  be  impure  than  a  swift  one, 
because  the  latter  tends  to  purify  itself  by  exposing  its 
impurities  to  the  oxidizing  power  of  the  air.  Ocean  water 
contains  a  large  proportion  of  common  salt.  The  propor- 
tions of  the  solid  substances  in  their  order  of  abundance 
are  shown  in  the  following  — 

TABLE  OF  SOLID  SUBSTANCES  IN  THE  OCEAN 


SUBSTANCE 

PER  CENT 

SUBSTANCE 

PER  CENT 

Sodium  Chloride  .  .  . 
Magnesium  Chloride  .  . 
Magnesium  Sulphate  .  . 
Calcium  Sulphate  .  . 

77.76 
10.88 
4.74 
3.60 

Potassium  Sulphate    .     . 
Calcium  Carbonate    .     . 
Magnesium  Bromide  .     . 
Other  Substances  .     .     . 

2.46 
.34 

.22 
traces 

The  peculiar  taste  of  ocean  water  is  due  to  the  presence 
of  these  substances,  and  since  the  water  only  is  removed 
by  evaporation,  the  ocean  always  has  a  "  salty  "  taste. 

Drinking  Water.  —  Water  used  as  a  beverage  should 
of  course  be  as  pure  as  possible.  As  a  rule  the  mineral 
matter  in  water  selected  for  drinking  is  not  injurious  to 
health;  but  since  water  may  become  contaminated  with 


54  INORGANIC  CHEMISTRY 

bacteria  which  produce  diseases  such  as  typhoid  fever  and 
cholera,  it  is  usually  necessary  to  purify  the  water  before  use. 

The  problem  of  obtaining  suitable  drinking  water  in  large 
quantities  is  local.  The  water  of  many  cities  is  purified 
by  filtering  it  through  a  layer  of  sand  and  gravel,  an  acre 
or  more  in  area  and  several  feet  deep.  Such  a  filter  removes 
bacteria  almost  completely,  though  it  must  be  frequently 
cleaned.  Sometimes  the  water  is  stored  in  a  large  settling 
basin  or  reservoir  and  purified  by  adding  alum,  or  a  similar 
substance,  which  causes  the  suspended-  matter  to  settle. 
Ozone  is  used  as  a  purifier  in  some  localities,  and  bleaching 
powder  has  been  applied  with  excellent  results  to  stored 
water  contaminated  with  certain  kinds  of  organic  matter. 
Dissolved  substances  cannot  be  removed  without  consid- 
erable difficulty,  so  as  a  rule  water  is  taken  from  a  source 
which  is  reasonably  pure. 

The  purity  of  drinking  water  is  usually  determined  by 
a  water  analysis.  This  is  not  a  decomposition  of  water, 
but  a  chemical  examination  of  a  sample  for  the  presence 
and  amount  of  certain  substances  which  indicate  or  cause 
impurity.  A  chemical  examination  is  of  limited  value, 
however,  unless  it  is  supplemented  by  a  microscopic  study 
of  a  fresh  sample  and  a  rigid  sanitary  inspection  of  the 
premises.  Water  which  is  clear,  sparkling,  cool,  attractive 
to  the  eye,  and  pleasant  to  the  taste  may  be  seriously  pol- 
luted by  disease  germs ;  or  it  may  be  liable  to  sudden  con- 
tamination from  some  unsuspected  source.  On  the  other 
hand,  a  rather  unpleasant-looking  water  may  be  harmless; 
hence  the  necessity  of  careful  and  extended  examination  of 
water  to  be  used  as  a  beverage. 

The  purification  of  water  may  be  readily  accomplished 
by  distillation.  This  operation  consists  essentially  in  boil- 
ing the  water,  condensing  the  resulting  vapor,  and  collecting 
the  liquid ;  by  this  method  the  non- volatile  matter  dissolved 


GENERAL  PROPERTIES  OF  WATER 


55 


in  the  water  remains  behind  in  the  distilling  vessel.  It  is 
performed  in  the  laboratory  by  means  of  a  condenser,  which 
is  shown  in  Figure  7.  The  condenser  consists  of  an  outer 
tube,  A  A,  provided  with  an  inlet  and  an  outlet  for  a  current 
of  cold  water,  which  surrounds  the  inner  tube,  BB.  The 
vapor  from  the  water  boiling  in  the  flask,  C,  condenses 
in  the  inner  tube,  owing  to  the  decrease  in  temperature, 


FIG.  7.  —  Condenser  arranged  for  the  distillation  of  water. 

and  drops  off  the  lower  end  of  this  tube,  as  the  distillate, 
into  the  receiver,  D,  while  the  non-volatile  impurities  remain 
behind  in  the  flask.  Distilled  water  is  prepared  on  a  large 
scale  in  metal  vessels,  and  the  vapor  is  condensed  in  a  block 
tin  pipe  coiled  inside  a  vessel  through  which  cold  water 
flows.  Distilled  water  is  used  in  the  chemical  laboratory; 
large  quantities  are  made  into  ice,  and  considerable  (espe- 
cially after  aeration)  is  used  as  a  beverage.  Distillation  is 
done  on  a  large  scale  by  boiling  water  in  a  metal  vessel 
and  condensing  the  vapor  in  a  spiral  tin  tube  cooled  by 
water. 


56  INORGANIC  CHEMISTRY 

Physical  Properties  of  Pure  Water.  —  Owing  to  its  marked 
solvent  power,  water  is  never  found  pure  in  nature,  and  is 
purified  even  in  the  laboratory  only  by  taking  special  pre- 
cautions. At  the  ordinary  temperature  pure  water  is  a 
tasteless  and  odorless  liquid.  It.  is  usually  colorless  and 
transparent,  but  thick  layers  are  bluish.  Water  is  a  poor 
conductor  of  heat.  This  last  property  can  be  shown  by 
boiling  water  near  the  surface  in  a  large  test  tube  containing 
a  piece  of  ice  weighted  down  upon  the  bottom.  The  ice 
remains  unmelted  for  some  time,  although  the  water  is 
boiling  a  few  inches  above  it. 

Most  liquids  expand  when  heated  and  contract  when 
cooled.  Water  behaves  exceptionally.  If  water  at  100°  C. 
is  gradually  cooled,  it  contracts  until  4°  C.  is  reached;  if 
the  cooling  continues,  the  water  expands  as  long  as  the  liquid 
state  is  maintained.  Hence  at  4°  C.  a  given  volume  contains 
the  greatest  weight  of  water.  In  other  words,  water  has 
its  maximum  density  at  4°  C.  The  density  of  water  at  4°  C. 
is  taken  as  1 ;  and  water  at  this  temperature  is  the  standard 
for  determining  the  specific  gravity  of  solids  and  liquids. 
Thus,  when  we  say  specific  gravity  of  gold  is  19,  we  mean 
that  a  piece  of  gold  is  19  times  heavier  than  an  equal  volume 
of  water  at  4°  C.  A  cubic  centimeter  of  water  at  4°  C. 
weighs  1  gm. 

The  expansion  of  water  when  cooled  from  4°  C.  to  0°  C. 
is  slight,  but  the  change  is  exceedingly  important  in  nature. 
When  the  water  on  the  surface  of  a  lake  or  river  cools,  it 
contracts,  and  since  it  is  heavier  (volume  for  volume)  than 
the  warmer  water  beneath,  it  sinks.  The  warmer  water 
rises,  becomes  cool,  and  likewise  sinks,  thus  causing  a  cir- 
culation which  continues  until  all  the  water  from  the  sur- 
face to  the  bottom  has  a  temperature  of  4°  C.  Now  if  the 
cooling  continues,  the  surface  water  expands  and  remains 
on  the  top,  because  it  is  lighter  than  the  water  beneath. 


GENERAL  PROPERTIES  OF  WATER      57 

Hence  when  the  temperature  of  the  air  falls  to  0°  C.,  this 
upper  layer  of  water  freezes  and  protects  the  remaining 
water  from  the  cold  air,  thus  stopping  the  circulation. 
Should  the  circulation  continue,  as  the  temperature  fell 
from  4°  C.  to  0°  C.,  the  whole  body  of  water  would  finally 
freeze  from  top  to  bottom.  This  condition  would  not  only 
destroy  aquatic  life,  but  profoundly  affect  climate. 

Water  solidifies  or  freezes  at  0°  C.  (or  32°  Fahrenheit). 
And  when  water  freezes,  it  expands  about  one  tenth  of  its  vol- 
ume. That  is,  100  cc.  of  water  produce  about  110  cc.  of  ice. 
In  other  words,  100  cc.  of  water  and  110  cc.  of  ice  weigh 
100  gm.  each.  Hence  ice  floats.  The  specific  gravity  of  ice  is 
about  .92.  The  pressure  exerted  by  water  when  it  freezes 
is  powerful.  Vessels  or  pipes  filled  with  water  often  burst 
when  the  water  freezes.  It  is  an  erroneous  but  popular 
idea  that  "  thawing  out  "  a  pipe  bursts  it.  As  a  matter 
of  fact,  ice  contracts  when  it  melts.  The  pipe  cracks  as 
soon  as  the  water  freezes,  and  when  the  ice  melts  a  channel 
is  left  for  the  water  to  flow  out  of  the  pipe.  Because  of  this 
property,  ice  is  an  effective  agent  in  splitting  rocks.  Water 
creeps  into  the  cracks,  especially  into  narrow  ones,  by  capil- 
lary attraction,  and  when  it  freezes,  the  rock  splits.  Water 
in  freezing  also  destroys  the  tissue  of  living  plants,  which 
are  often  said  to  have  been  "touched  by  frost."  Frozen 
flesh  for  a  similar  reason  becomes  pulpy  and  is  more  liable 
to  putrefy  when  thawed. 

Ice  melts  at  0°  C.  (32°  F.),  which  is  also  the  freezing  point 
of  water.  Ice  often  crystallizes  in  freezing,  but  the  individ- 
ual crystals  are  seldom  visible  except  during  the  first  stages 
of  the  process.  Snow  crystals  are  common.  They  are 
always  six-sided,  and  are  formed  in  the  atmosphere  by  the 
freezing  of  water  vapor. 

Water  evaporates  at  all  temperatures,  passing  off  as  an 
invisible  vapor  into  the  atmosphere  or  into  the  air  confined 


58 


INORGANIC    CHEMISTRY 


over  it.  If  water  is  heated  in  an  open  vessel,  the  tem- 
perature rises  and  vapor  passes  off  rapidly  until  the  ther- 
mometer reaches  100°  C.  (or  212°  F.).  At  this  point  water 
boils ;  i.e.  it  changes  rapidly  into  vapor  without  rise  of  tem- 
perature. This  vapor,  if  allowed  to  escape,  cools  and  con- 
denses quickly  into  a  cloud  of  minute  drops  of  water.  This 
cloud  is  popularly  called  steam.  Accu- 
rately speaking,  steam  is  invisible.  What  I  U  \\B  1 1C 
we  call  steam  is  a  cloud  or  collection 


FIG.  8.  —  Vapor  pressure. 

of  very  small  particles  of  water.  This  may  be  illustrated 
by  boiling  water  in  a  large  glass  flask.  The  inside  of  the 
flask  is  perfectly  transparent,  although  there  is  a  cloud  of 
"  steam  "  issuing  from  its  mouth. 

Escaping  vapor  exerts  pressure,  as  may  be  readily  shown. 
If  a  little  water  is  introduced  into  a  dry  closed  bottle  having 
a  U-shaped  tube  of  colored  liquid  connected  with  the  interior 
to  serve  as  a  gauge,  the  difference  in  the  levels  of  the  colored 
liquid  indicates  a  pressure  inside  the  bottle  (Fig.  8).  This 
pressure  is  due  to  vapor  escaping  from  the  water.  The 
ability  of  water  (or  of  any  other  volatile  liquid)  to  generate 
vapor  is  called  the  vapor  tension  of  the  liquid.  The  pressure 
exerted  by  the  vapor  is  called  the  vapor  pressure  of  water. 


GENERAL  PROPERTIES  OF  WATER  59 

The  amount  of  vapor  pressure  depends  on  the  temperature. 
This  is  seen  by  comparing  the  heights  of  the  mercury  in  the 
barometer  tubes  in  Figure  8  (right) .  In  the  tube  A  there  is 
no  water  vapor  in  the  space  above  the  mercury,  and  there- 
fore the  height  of  the  mercury  is  760  mm.  In  the  tube  B 
the  space  above  the  mercury  is  filled  with  water  vapor  at  20° 
C. ;  the  vapor  exerts  a  pressure  and  forces  the  mercury  down 
to  nearly  742  mm.  That  is,  the  water  vapor  at  20°  C.  exerts 
a  pressure  equal  to  about  18  mm.  of  mercury.  Similarly,  in 
the  tube  C  the  space  is  filled  with  water  vapor  at  50°  C.  and 
the  mercury  is  forced  down  to  678  mm.,  the  water  vapor 
exerting  a  pressure  of  about  82  mm.  If  the  vapor  were  at 
0°  C.,  the  vapor  pressure  would  be  about  4.5  mm.,  and  at 
100°  C.  the  vapor  pressure  would  be  760  mm.  The  latter 
value  is  instructive,  for  it  means  that  at  the  boiling  point  of 
water  (100°  C.)  the  vapor  pressure  just  balances  the  normal 
atmospheric  pressure.  The  pressure  exerted  by  water  vapor, 
as  shown  in  Figure  8,  is  independent  of  atmospheric  pressure, 
the  size  and  shape  of  the  inclosing  space,  and  the  pressure 
of  other  gases.  It  depends  solely  on  the  temperature  of 
the  evaporating  water,  and  has  a  maximum  value  for  each 
temperature.  The  maximum  value  represents  the  vapor 
tension  of  water  at  the  given  temperature  (see  Equilibrium, 
page  61).  These  values  can  be  found  in  the  App.,  §  4. 

A  practical  application  of  vapor  pressure  is  made  in  de- 
termining the  weight  of  a  liter  of  oxygen  and  in  similar 
experiments  where  gases  are  measured  over  water.  The 
oxygen  gas  is  collected  in  a  bottle  or  graduated  tube  inverted 
in  a  vessel  of  water.  If  the  gas  is  allowed  to  stand  confined 
over  the  water  long  enough,  it  becomes  saturated  with  water 
vapor ;  i.e.  the  tube  finally  contains  a  mixture  of  oxygen 
and  the  maximum  amount  of  water  vapor  at  the  given  tem- 
perature. In  such  a  mixture,  where  no  chemical  action 
occurs,  each  gaseous  constituent  shares  the  total  pressure 


60  INORGANIC  CHEMISTRY 

(against  the  atmospheric  pressure).  This  proportionate 
part  of  the  total  pressure  is  called  the  partial  pressure  of 
that  gas.  If  the  oxygen  were  confined  over  mercury,  its 
pressure  would  be  the  same  as  its  partial  pressure  when 
saturated  with  water  vapor.  Hence  the  actual  pressure 
under  which  the  oxygen  itself  exists  is  found  by  determining 
its  partial  pressure  or  by  subtracting  the  partial  pressure 
of  the  water  vapor  from  the  total  pressure  (indicated 
by  the  barometer).  The  latter  method  is  used,  because  the 
pressure  of  water  vapor  at  any  temperature  is  known  and 
can  be  taken  directly  from  the  table.  Incorporating  this 
fact  into  the  formula  given  in  Chapter  IV  for  reducing 
the  volume  of  a  gas  to  its  volume  at  0°  C.  and  760  mm.  the 
formula  becomes  — 

V'(P'-a) 
"760  [1  +  (.00366  0] ' 

In  this  formula  V  means  the  volume  of  dry  oxygen  at  0°  C. 
and  760  mm.,  and  a  means  the  vapor  pressure  (found  in  the 
table  in  the  App.,  §  4).  (See  also  Laboratory  Manual,  App.  B.) 
Several  important  conclusions  can  be  drawn  from  the 
discussion  of  vapor  pressure  in  the  foregoing  paragraphs. 
Recalling  the  fact  that  the  vapor  pressure  of  water  at  100°  C. 
is  760  mm.,  it  is  obvious  that  this  temperature  is  the  boiling 
point  at  this  pressure.  Water  boils  when  its  vapor  escapes 
with  sufficient  pressure  to  overcome  the  pressure  upon  its 
surface.  Hence  the  boiling  point  depends  upon  the  pressure 
—  either  of  the  atmosphere  or  of  the  vapor  in  the  vessel. 
The  boiling  point  of  water  is  100°  C.  (or  212°  F.)  only  when 
the  atmospheric  pressure  is  normal,  i.e.  760  mm.  The  boil- 
ing point  becomes  lower  as  the  pressure  is  decreased  and 
higher  as  the  pressure  is  increased.  For  example,  in  the 
city  of  Mexico  (7500  ft.  above  the  sea  level)  water  boils  at 
92°  C.,  and  in  Quito  (9350  ft.  above  the  sea  level)  water 


GENERAL  PROPERTIES  OF  WATER      61 

boils  at  about  90°  C.  Again,  the  condition  in  a  closed 
vessel  containing  water  and  water  vapor  merits  considera- 
tion. As  the  water  evaporates,  the  vapor  pressure  increases 
until  the  space  above  the  water  becomes  saturated  with 
water  vapor.  Then  the  vapor  pressure  remains  constant  as 
as  long  as  the  temperature  is  fixed,  say  20°  C.  The  vapor 
tension  of  the  water  is  now  equal  to  the  vapor  pressure  of  the 
vapor.  In  other  words,  vapor  tension  and  vapor  pressure 
balance  each  other.  Vapor  is  condensing,  however,  just  as 
fast  as  water  is  changing  into  vapor.  Such  a  condition  of 
mutual  exchange  is  called  a  state  of  equilibrium,  i.e.  a  con- 
dition in  which  two  opposing  processes  balance  but  do  not 
stop  each  other.  Equilibrium  is  approached  or  reached  in 
many  processes,  chemical  as  well  as  physical. 

Chemical  Properties  of  Water.  —  Water  has  such  con- 
spicuous physical  properties  that  its  chemical  properties 
are  sometimes  ignored  or  discussed  in  connection  with  other 
substances.  Such  discussion,  it  must  be  admitted,  is  often 
more  appropriate  elsewhere,  but  certain  phases  of  this 
topic  need  attention  here.  Water  at  the  ordinary  tempera- 
ture interacts  with  certain  metals,  especially  sodium  and 
potassium.  This  chemical  change,  as  already  stated,  is 
substitution.  •  The  metal  is  substituted  for  part  of  the 
hydrogen  of  the  water,  thereby  liberating  hydrogen  and 
producing  a  new  compound,  which  contains  sodium,  hydro- 
gen, and  oxygen,  and  is  called  sodium  hydroxide.  Mag- 
nesium and  zinc  interact  similarly  with  boiling  water,  the 
products  being  hydrogen,  magnesium  hydroxide,  and  zinc 
hydroxide.  Water  is  decomposed  to  some  extent  into  its 
component  elements  (oxygen  and  hydrogen)  by  intense 
heat;  at  about  2000°  C.  the  decomposition  is  less  than  2 
per  cent.  As  the  temperature  falls,  the  elements  recombine 
to  form  water.  Both  decomposition  and  recombination 


62  INORGANIC  CHEMISTRY 

are  gradual.  Water  when  decomposed  by  heat  behaves 
differently  from  potassium  chlorate.  The  latter  compound 
suffers  permanent  decomposition;  i.e.  the  oxygen  and  the 
potassium  chloride  into  which  it  is  decomposed  do  not  re- 
combine  as  the  temperature  falls.  It  is  customary  to 
distinguish  these  two  kinds  of  decomposition.  That  kind 
which  takes  place  gradually  at  high  temperatures  and  is 
followed  by  recombination  when  the  temperature  is  lowered 
is  called  dissociation  by  heat.  Several  common  compounds 
undergo  dissociation  when  heated,  and  they  will  be  discussed 
later.  Water  is  decomposed  when  heated  with  certain 
metals,  e.g.  iron,  but  the  oxygen  at  once  combines  with 
the  metal,  so  the  final  products  are  hydrogen  and  an  oxide 
of  the  metal.  A  mixture  of  water  and  sulphuric  acid  yields 
hydrogen  and  oxygen  when  subjected  to  the  action  of  an 
electric  current.  Apparently  the  water  is  merely  decom- 
posed into  its  elements,  but  it  will  be  shown  subsequently 
that  the  action  is  not  so  simple.  (See  Chapter  IX.)  Water 
combines  directly  with  many  oxides.  Thus,  lime,  which 
has  the  chemical  name  calcium  oxide,  combines  directly 
with  water  and  forms  a  compound  called  calcium  hydroxide. 
Similarly,  barium  oxide  forms  barium  hydroxide,  sulphur 
dioxide  forms  sulphurous  acid,  phosphorus  pentoxide  forms 
phosphoric  acid.  These  chemical  changes  may  be  typically 
represented  thus  :  — 

Calcium  Oxide  +  Water  =  Calcium  Hydroxide. 
Sulphur  Dioxide  +  Water  =  Sulphurous  Acid. 

Oxides  which  react  thus  with  water  are  called  anhydrides. 
(See  page  162.)  Water  also  combines  with  certain  solids 
when  they  separate  from  a  solution  by  crystallization. 
Thus,  copper  sulphate  crystals  are  blue,  but  when  heated, 
water  is  liberated  and  the  crystals  crumble  to  a  gray  white 
powder.  This  class  of  compounds,  in  which  water  is  a 


GENERAL  PROPERTIES  OF  WATER      63 

definite  component,  is  treated  below.     (See  in  this  chapter 
Solution  and  Crystallization.) 

Solvent  Power  of  Water.  —  This  is  one  of  the  most  con- 
spicuous properties  of  water  and  can  be  discussed  from 
several  standpoints.  Only  the  simpler  physical  aspects  are 
treated  in  this  chapter.  More  extended  discussion  may  be 
found  in  Chapter  IX. 

Daily  experience  shows  that  many  solids,  liquids,  and 
gases  disappear  when  put  into  water.  This  operation  is 
called  dissolving  or  putting  into  solution.  The  clear, 
transparent  liquid  containing  the  dissolved  substance  is 
called  the  solution.  The  liquid  in  which  the  substance 
dissolves  is  called  the  solvent,  and  the  dissolved  substance 
is  called  the  solute.  If  the  solute  is  not  volatile,  or  not  very 
volatile,  it  can  be  recovered  by  evaporating,  or  distilling  off, 
the  water.  The  degree  of  solubility  of  a  substance  is  often 
conveniently  expressed  by  the  terms  slightly  soluble,  soluble, 
and  very  soluble.  It  is  more  accurate,  however,  to  state 
the  proportions  of  the  solvent  and  the  solute,  and  also  the 
temperature.  Thus,  instead  of  saying  that  common  salt  is 
very  soluble  in  cold  water,  it  is  better  to  state  that  36  gm. 
of  salt  dissolve  in  100  gm.  of  water  at  20°  C.  There  is  a 
limit  to  the  weight  of  each  substance  which  a  given  weight 
of  water  will  dissolve.  That  is,  substances  differ  widely 
in  the  degree  of  solubility.  Some,  like  potassium  perman- 
ganate, are  very  soluble,  while  others,  like  sand,  dissolve 
only  very  slightly.  The  latter  class  is  often  said  to  be  in- 
soluble. Strictly  speaking,  no  substance  is  insoluble,  but 
the  term  is  often  applied  to  those  substances  whose  solu- 
bility is  so  slight  that  it  can  be  neglected  in  most  cases. 
Many  minerals  and  rocks  belong  to  ,the  class  of  so-called 
insoluble  substances. 

The  concentration  of  a  solution  is  its  strength  as  de- 


64  INORGANIC   CHEMISTRY 

ter mined  by  the  amount  of  solute  dissolved  in  a  given 
amount  of  solvent.  The  ratio  of  these  two  amounts  ex- 
presses the  concentration.  A  solution  which  contains  a 
small  proportion  of  solute  is  called  a  dilute  solution,  or  one 
of  small  concentration;  one  containing  a  large  proportion 
of  solute  is  called  a  concentrated  solution.  Thus,  dilute 
sulphuric  acid  usually  contains  one  part  of  acid  to  three 
or  more  parts  of  water,  while  concentrated  sulphuric  acid 
is  nearly  98  per  cent  acid.  Sometimes  the  terms  weak 
and  strong  are  used  instead  of  dilute  and  concentrated. 
Other  terms  are  defined  below.  (See  Solubility"  of  Solids.) 

Solution  of  Gases.  —  Water  dissolves  many  gases.  The 
solubility 'varies  widely.  For  example,  one  volume  of  water 
dissolves  about  1150  volumes  of  ammonia  gas  (at  0°  C. 
and  760  mm.),  550  of  hydrochloric  acid  gas,  80  of  sulphur 
dioxide  gas,  1.8  of  carbon  dioxide  gas,  .04  of  oxygen,  .02  of 
nitrogen,  and  .02  of  hydrogen. 

As  a  rule  the  volume  of  gas  dissolved  by  a  given  volume 
of  water  decreases  with  rise  of  temperature.  Thus,  when 
ammonium  hydroxide  is  heated,  ammonia  gas  escapes 
freely.  So  also,  when  water  is  heated  gradually,  the  air 
dissolved  in  the  water  gathers  as  little  bubbles,  which  soon 
rise  and  escape.  Again,  100  volumes  of  water  dissolve 
approximately  4  volumes  of  oxygen  at  0°  C.,  3  at  20°  C., 
1.8  at  50°  C.,  and  none  at  100°  C.  The  complete  removal 
of  a  gas  by  boiling  its  solution  is  not  possible  in  the  case 
of  certain  very  soluble  gases,  like  hydrochloric  acid  gas. 
Such  solutions  when  heated  lose  gas  or  water  until  a  certain 
concentration  is  reached,  and  then  the  solution  boils  at  a 
constant  temperature  (see  pages  206,  218). 

Pressure  influences  the  solubility  of  gases.  Thus,  large 
quantities  of  carbon  dioxide  gas  are  forced  into  cylinders 
full  of  water  in  preparing  soda  water.  When  the  pressure 


GENERAL  PROPERTIES   OF  WATER  65 

is  removed,  the  gas  escapes  rapidly  and  causes  the  soda 
water  to  froth  or  foam.  This  rapid  escape  of  gas  is  called 
effervescence.  Underground  waters  often  contain  large 
amounts  of  gases,  especially  carbon  dioxide,  owing  to  the 
great  pressure  to  which  subterranean  gases  are  subjected. 
Hence,  mineral  waters  often  effervesce  when  they  come  to 
the  surface.  The  greater  the  pressure,  the  greater  the 
amount  of  gas  dissolved.  More  accurately  stated,  the  weight 
of  a  moderately  soluble  gas  dissolved  by  a  given  weight  of 
water  is  directly  proportional  to  the  pressure,  if  the  tem- 
perature is  constant.  This  is  one  form  of  the  general  state- 
ment known  as  Henry's  Law.  Gases  which  are  very  soluble 
or  which  dissolve  under  great  pressure  deviate  from  this 
law,  probably  owing  to  chemical  combination.  (See  also 
preceding  paragraph.) 

When  a  mixture  of  gases,  such  as  air,  dissolves  in  water, 
each  ingredient  behaves  independently  of  the  other.  More 
strictly,  each  gas  dissolves  proportionally  to  its  partial 
pressure.  For  example,  the  ratio  of  oxygen  to  nitrogen  in 
air  is  about  1 : 4  and  the  ratio  of  their  solubilities  when  not 
mixed  is  2:1,  but  the  ratio  of  oxygen  to  nitrogen  in  water 
saturated  with  air  (at  760  mm.)  is  7 : 13  or  about  1:2.  The 
difference  is  due  to  the  fact  that  their  solubility  when  mixed 
is  determined  not  merely  by  their  power  to  dissolve  as  free 
gases  but  also  by  their  partial  pressure. 

Solutions  of  Liquids.  — The  solubility  of  liquids  in  water 
varies  between  wide  limits.  Some,  such  as  alcohol  and 
glycerin,  are  soluble  in  all  proportions.  Oils,  such  as 
kerosene,  are  practically  insoluble;  hence  the  old  adage, 
"  Oil  and  water  will  not  mix."  Carbon  disulphide  is  also 
almost  entirely  insoluble,  as  is  shown  by  the  fact  that  after 
agitation  with  water  it  separates  almost  entirely  as  a  distinct 
layer  ;  being  heavier  than  water,  this  layer  forms  at  the 


66  INORGANIC  CHEMISTRY 

bottom.  The  mere  formation  in  this  way  of  separate  layers 
by  two  liquids  is  not  conclusive  evidence  of  relative  insolu- 
bility. Only  in  those  cases  which  exhibit  perfect  mutual 
solubility  is  the  separation  into  layers  after  agitation  im- 
possible. Ether  and  water  form  two  layers,  but  each  dis- 
solves appreciably  in  the  other.  Thus,  about  2  gm.  of 
ether  dissolves  in  100  gm.  of  water  (at  20°  C.)  and  about 
10  gm.  of  water  dissolves  in  100  gm.  of  ether.  The  upper 
layer  consists  of  ether  saturated  with  water ;  the  lower, 
of  water  saturated  with  ether.  Alcohol  and  water  form  no 
such  layers,  not  because  each  is  soluble  in  the  other,  but 
because  each  is  soluble  without  limit  in  the  other  ;  i.e.  it  is 
a  case  of  perfect  mutual  solubility.  In  many  cases  a  rise  in 
temperature  increases  the  solubility  of  liquids  in  water. 

Solutions  of  Solids.  —  The  solubility  of  solids  in  water  is 
a  subject  of  much  practical  importance.  The  abundance 
of  water  and  its  power  to  dissolve  such  a  vast  number  of 
different  solids  have  led  some  to  call  water  "  the  universal 
solvent."  The  far-reaching  effect  of  this  marvelous  power 
in  nature  and  the  indispensable  value  of  water  to  man  have 
been  considered.  (See  above.) 

The  degree  of  solubility  of  solids  in  water  depends  upon 
the  substance  itself  and  the  temperature  of  the  water.  Some 
solids  are  very  soluble,  while  others  are  difficultly  soluble. 
In  most  cases  solubility  increases  with  a  rise  of  tempera- 
ture ;  hence  the  common  practice  of  heating  to  hasten  solu- 
tion. The  effect  of  an  increase  of  temperature  on  solubility 
is  sometimes  very  marked,  the  solubility  being  increased 
many  fold  in  passing  from  the  ordinary  temperature  to  the 
boiling  point.  A  few  solids  (e.g.  calcium  hydroxide)  are  less 
soluble  in  hot  water  than  in  cold,  and  a  few  others  (e.g.  sodium 
chloride)  dissolve  to  about  the  same  degree  in  hot  and  cold 
water.  These  properties  of  the  solutions  of  various  solids 
are  illustrated  by  the  following  — 


GENERAL   PROPERTIES   OF  WATER  67 

TAKLE  OF  THE  SOLUBILITY  OF  SOLIDS  IN  WATER 


SOLIDS 

NUMBER  OF  GRAMS  IN  SOLUTION  IN 
100  GRAMS  OF  WATER 

20°  C. 

100"?  C. 

Calcium  Chloride    .         .... 

74.5 
42.3 
36.2 
7.2 
35 
13 
31.6 
10.6 
36 

159 
203.3 
73.8 
65.9 
57 
.102 
246 
26 
39.8 

Potassium  Chlorate     

Potassium  Chloride     

Potassium  Nitrate  

Potassium  Sulphate     . 

Sodium  Chloride     

Other  facts  about  the  solubility  of  solids  are  illustrated  by 
this  table  besides  the  dependence  on  the  substance  itself 
and  on  temperature,  Inspection  of  the  table  shows  that 
there  is  a  limit  to  the  solubility  of  solids  in  water  at  a  given 
temperature.  That  is,  a  given  weight  of  water  at  a  fixed 
temperature  will  dissolve  a  definite  weight  of  solid  and  no 
more,  even  though  some  undissolved  solid  is  .in  the  liquid. 
A  solution  conforming  to  the  conditions  just  stated  is  said 
to  be  saturated  or  to  have  reached  its  maximum  concentra- 
tion. Thus,  100  gm.  of  water  holds  7.2  gm.  of  potassium 
chlorate  in  solution  at  20°  C.  If  more  potassium  chlorate  is 
added,  it  remains  undissolved  (provided  the  temperature  of 
the  solution  and  the  weight  of  water  are  unchanged).  If 
the  temperature  falls,  more  solid  comes  out  of  solution ; 
and  conversely,  if  the  temperature  rises,  more  solid  dissolves. 
As  long  as  the  maximum  concentration  is  maintained  in 
contact  with  some  of  the  undissolved  solid,  the  solution  is 
saturated.  The  ratio  of  the  weight  of  the  solute  to  the  weight 
of  the  solvent  in  a  saturated  solution  is  called  the  solubility 
of  the  solid.  Solubility  is  expressed  in  several  ways.  One 


68 


INORGANIC  CHEMISTRY 


way  represents  the  solvent  by  100  gm. ;  on  this  basis  the 
solubility  of  a  solid  becomes  the  number  of  grams  of  solid 
dissolved  by  100  gm.  of  water.  In  the  case  just  cited,  the 
maximum  concentration,  or,  as  it  is  more  often  called,  the 


10°    20°    so 


40°       50°      60°      70°      80J      90°      100° 
Temperature 
FIG.  9.  —  Solubility  curves. 


solubility,  of  potassium  chlorate  is  7.2  gm.  at  20°  C.  and 
55.9  gm.  at  100°  C. 

The  table  of  solubilities  just  given  is  limited  to  two  tem- 
peratures.    A  convenient  way  of  showing  the  solubility  of 


GENERAL  PROPERTIES  OF  WATER      69 

a  substance  as  the  temperature  varies  between  convenient 
points  is  by  a  solubility  curve.  The  curves  of  several  sub- 
stances are  shown  in  Figure  9.  The  temperatures  are  read 
from  the  vertical  lines  and  the  number  of  grams  of  solute  in 
100  gm.  of  water  is  read  from  the  horizontal  lines.  For 
example,  if  we  wish  to  know  the  temperature  at  which 
40  gm.  of  potassium  chlorate  are  held  in  solution  by  100  gm. 
of  water,  it  is  only  necessary  to  find  where  the  horizontal 
line  numbered  40  cuts  the  potassium  chlorate  curve,  and 
then  follow  the  vertical  line  down  to  the  temperature  num- 
bers, where  80°  C.  is  found.  Similarly,  100  gm.  of  water 
dissolve  37  gm.  of  sodium  chloride  at  60°  C.,  while  the  same 
weight  of  water  dissolves  110  gm.  of  potassium  nitrate  at 
60°  C.,  and  so  on. 

When  hot  solutions  are  cooled  or  concentrated  solutions 
are  evaporated,  the  solute  separates  in  the  solid  state  just 
as  soon  as  the  saturation  point  (at  a  lower  temperature)  is 
passed.  Under  certain  favorable  conditions  the  solid  is 
deposited  in  masses  having  a  definite  form.  These  masses 
are  called  crystals  and  the  process  of  obtaining  them  is 
called  crystallization.  The  form  and  color  of  the  crystals 
are  characteristic  of  the  particular  substance  and  serve  to 
identify  it.  Thus,  potassium  chlorate  crystallizes  in  shin- 
ing white  plates  or  leaves,  and  common  salt  in  white  cubes. 
The  deposition  of  crystals  is  not  always  as  prompt  as  just 
stated.  Thus  a  hot,  concentrated  solution  of  some  solids, 
such  as  sodium  acetate  and  sodium  thiosulphate,  deposits  no 
crystals  even  when  the  clear  solution  cools.  Such  solutions 
are  called  supersaturated.  Supersaturation  can  occur  only 
when  the  undissolved  solid  is  not  present.  If  a  fragment  of 
the  solid  is  dropped  into  the  supersaturated  solution,  crystals 
very  soon  begin  to  form  upon  the  fragment,  and  this  separa- 
tion continues  until  just  enough  solid  is  left  in  solution  to 
produce  saturation  at  the  prevailing  temperature.  The 


70  INORGANIC   CHEMISTRY 

amount    of    solid    thus    separated   is  often  very  great  and 
sometimes  forms  a  solid  mass  in  the  test  tube. 

It  is  evident  from  the  foregoing  paragraphs  that  there  are 
three  general  classes  of  solutions  of  solids,  viz.  unsaturated, 
saturated,  and  supersaturated.  These  classes  can  be  dis- 
tinguished by  bringing  each  in  contact  with  more  of  the 
solid.  If  the  solution  is  unsaturated,  more  solid  will  dis- 
solve; if  saturated,  no  more  will  dissolve;  if  supersaturated, 
solid  will  be  deposited  until  saturation  is  reached.  In  an 
unsaturated  solution  the  concentration  is  less  than  in  a 
saturated  solution,  while  in  a  supersaturated  solution  it  is 
greater.  If  we  think  of  these  solutions  as  being  in  contact 
with  a  solid,  their  relations  will  be  clearer  than  if  we  regard 
them  as  reservoirs,  so  to  speak,  for  more  or  less  solid.  In  the 
saturated  solution  there  is  equilibrium  between  the  solution 
and  the  solid.  That  is,  there  is  an  equal  tendency  for  the 
solid  to  enter  and  to  leave  the  solution.  But  in  the  other 
two  solutions  no  such  equilibrium  prevails,  for  the  un- 
saturated solution  takes  up  more  solid  and  the  supersatu- 
rated solution  deposits  solid;  both  are  stable  when  solid  is 
absent,  but  unstable  as  soon  as  solid  is  present. 

Solution  and  Crystallization.  — Under  the  chemical  prop- 
erties of  water  it  was  stated  that  water  combines  with  cer- 
tain solids  when  they  are  separated  from  a  solution  by 
crystallization.  Crystals  deposited  from  solutions  often 
contain  water,  which  is  an  essential  part  of  the  compound. 
The  combined  water  must  not  be  confused  with  water  which 
adheres  to  a  crystal  or  is  inclosed  in  it.  Even  after  the  crys- 
tals are  powdered  and  dried,  the  combined  water  remains, 
which  can  be  removed  by  heat  or  sometimes  merely  by  ex- 
posure to  the  air.  Loss  of  water  is  sometimes  attended  by 
loss  of  color  and  always  by  loss  of  crystalline  appearance. 
Thus,  crystallized  sodium  carbonate  turns  dull  and  crumbles 


GENERAL   PROPERTIES   OF   WATER  71 

in  the  air;  blue  crystallized  copper  sulphate  turns  white 
slowly  at  the  ordinary  temperature  and  very  rapidly  when 
heated,  finally  becoming  a  gray  powder;  but  the  variety  of 
crystallized  gypsum  called  selenite  must  be  heated  before 
the  combined  water  passes  off,  whereupon  the  crystal  be- 
comes chalky  and  crumbles  when  compressed.  The  pro- 
portion of  combined  water  in  crystals  is  not  arbitrary.  It  is 
constant  in  the  same  compound  when  crystallized  .under 
uniform  conditions.  The  amount  in  different  substances 
varies  between  wide  limits,  as  can  be  seen  by  the  following  — 

TABLE  OF  COMBINED  WATER  IN  CRYSTALS 


CRYSTALLIZED  SOLID 

PER  CENT  OF  COMBINED  WATER 

14.75 

Copper  Sulphate      

36.36 

Iron  Sulphate      . 

45.35 

There  is  no  doubt  that  the  water  in  the  crystallized  form 
of  compounds  is  not  chemically  combined  in  the  way  the 
other  elements  are ;  for  when  such  crystals  are  heated  with 
proper  precautions,  only  the  water  is  removed,  the  other 
constituents  of  the  original  compound  remaining  intact. 
In  most  cases  gentle  heating  suffices  to  expel  the  water. 
Hence  the  combination  between  the  water  and  the  rest  of 
the  compound  must  be  weaker  than  that  between  the  con- 
stituents of  the  residue  or  at  least  of  a  different  order.  This 
fact  is  often  indicated  by  separating  the  water  in  the  formula ; 
e.g.  CuSO4 .  5  H2O  and  BaCl2 .  2  H2O.  Water  chemically  com- 
bined in  a  crystal  and  readily  removed  in  a  definite  propor- 
tion by  heating  is  called  water  of  crystallization.  Compounds 
containing  water  of  crystallization  are  sometimes  called 
hydrates  or  hydrated  compounds.  Conversely,  compounds 
which  have  been  deprived  of  water  of  crystallization  are  said 


72  INORGANIC   CHEMISTRY 

to  be  anhydrous  or  dehydrated.  For  example,  blue  crystal- 
lized copper  sulphate  is  a  hydrate  of  the  compound  copper 
sulphate  CuSO4,  but  when  the  blue  compound  is  heated,  it 
becomes  anhydrous  or  dehydrated  copper  sulphate,  which 
is  a  gray  powder.  Anhydrous  compounds  often  readily  be- 
come hydrated  again.  Thus,  when  the  gray  anhydrous 
copper  sulphate  is  added  to  water,  a  blue  solution  is  pro- 
duced .from  which  blue  crystals  of  hydrated  copper  sulphate 
are  readily  obtained. 

Some  substances  give  up  their  water  of  crystallization 
wholly  or  in  part  upon  mere  exposure  to  the  air;  such  sub- 
stances are  said  to  be  efflorescent  or  to  effloresce.  Thus, 
sodium  carbonate  crystals  when  left  in  the  air  become 
opaque  and  ultimately  crumble,  owing  to  the  slow  escape 
of  their  water  of  crystallization.  The  white  spots  often  seen 
on  blue  crystals  of  copper  sulphate  or  on  green  ones  of 
iron  sulphate  are  due  to  efflorescence.  An  explanation  of 
efflorescence  is  found  in  the  principle  of  vapor  pressure. 
Substances  containing  water  of  crystallization  exert  a  vapor 
pressure.  If  this  vapor  pressure  is  greater  than  the  pressure 
of  the  water  vapor  in  the  atmosphere,  the  substance  loses 
water  until  the  vapor  pressures  are  equal  or  until  all  the 
water  has  escaped  from  the  substance.  Hence,  in  general, 
all  hydrated  compounds  effloresce,  if  they  exert  a  vapor 
pressure  greater  than  that  of  the  atmosphere  (at  that  time). 

It  should  not  be  concluded  from  the  foregoing  statements 
that  all  crystallized  solids  which  dissolve  in  water  contain 
water  of  crystallization.  Many  do  not ;  e.g.  sodium  chloride, 
potassium  chlorate,  sugar,  potassium  nitrate,  and  potassium 
dichromate.  No  satisfactory  explanation  has  been  given 
for  the  absence  of  water  of  crystallization  in  certain  crystal- 
lized compounds;  nor  for  the  varying  amount  and  relation 
to  color  and  crystal  form  in  hydrated  compounds. 

The  formation  of  crystals  is  not  limited  to  a  single  process, 


GENERAL  PROPERTIES  OF  WATER      73 

viz.  cooling  or  evaporating  a  solution.  It  can  be  accom- 
plished by  fusion  and  by  sublimation.  Thus,  sulphur 
crystallizes  when  melted  (or  fused)  and  then  cooled,  and 
iodine  crystallizes  when  vaporized  and  then  cooled.  These 
processes  of  obtaining  crystals  are  called  respectively  evapo- 
ration, fusion,  and  sublimation.  All  three  will  be  illustrated 
in  succeeding  chapters.  As  a  rule  crystals  obtained  by  any 
of  these  methods  are  quite  pure,  and  crystallization  is  one 
of  the  processes  frequently  used  in  the  industrial  preparation 
of  chemicals.  A  crystal  is  "a  solid  body  bounded  by  plane 
surfaces  arranged  according  to  definite  laws,  and  possessing 
definite  physical  properties;  both  external  form  and  physical 
properties  resulting  from,  and  being  the  expression  of,  definite 
internal  structure."  The  external  form  is  the  most  con- 
spicuous characteristic  of  a  crystal,  and,  as  a  rule,  each  sub- 
stance has  a  crystal  form  or  series  of  closely  related  forms 
by  which  it  can  be  distinguished.  Thus,  salt  crystallizes  in 
cubes,  alum  in  octahedrons,  sulphur  in  orthorhombic  forms, 
and  calcite  (calcium  carbonate)  in  hexagonal  forms.  Many 
minerals  occur  as  crystals,  and  since  they  are  the  natural 
form  of  chemical  compounds  as  well  as  the  source  of  many 
elements,  a  knowledge  of  the  common  crystal  forms  and 
the  properties  of  crystals  is  indispensable  in  identifying  sub- 
stances and  interpreting  descriptions.  Such  knowledge  is 
best  acquired  by  a  constant  examination  of  crystals.  The 
preliminary  treatment  given  in  the  Appendix,  §  3,  will  serve 
as  an  introduction  to  this  important  subject. 

Solution  and  Vapor  Pressure.  —  Aqueous  solutions  have 
a  vapor  pressure,  but  it  is  less  than  the  vapor  pressure  of 
water  at  the  same  temperature.  This  fact  can  be  illustrated 
by  introducing  a  solution  into  the  barometer  tube  as  described 
under  vapor  pressure.  Under  parallel  conditions  the  mer- 
cury will  always  stand  higher  in  the  tube  containing  the 


74  INORGANIC  CHEMISTRY 

solution,  i.e.  the  vapor  pressure  is  less,  and  consequently  the 
mercury  will  be  less  depressed.  Several  conclusions  can  be 
drawn  from  this  fact.  Since  the  boiling  point  depends  upon 
pressure,  solutions  must  be  heated  to  a  higher  temperature 
than  water  before  boiling  occurs.  That  is,  a  solid  dissolved 
in  water  elevates  the  boiling  point  of  water.  Moreover,  the 
elevation  of  the  boiling  point  depends  in  general  upon  the 
weight  of  solid  dissolved  in  a  given  weight  of  water  ;  i.e.  it 
is  proportional  to  the  concentration  of  the  solution.  This 
means  that  if  the  weight  of  solute  is  doubled,  the  elevation 
of  the  boiling  point  is  doubled.  Similar  statements  can  be 
made  about  the  freezing  point  of  solutions.  The  freezing 
point  of  a  solution  is  lower  than  the  freezing  point  of  water, 
and  the  depression  of  the  freezing  point  is  proportional  to 
the  concentration  of  the  solution.  Important  deductions 
will  subsequently  be  made  from  these  relations.  In  pass- 
ing, it  is  interesting  to  note  two  common  illustrations  of  the 
foregoing  statements,  viz.  that  boiler  water  containing  much 
dissolved  solid  has  to  be  heated  to  a  higher  temperature  than 
pure  water  in  the  production  of  steam,  and  that  the  salt 
water  along  the  seashore  freezes,  if  at  all,  with  more  diffi- 
culty than  the  fresh  water  of  near-by  rivers.  Finally,  the 
relatively  lower  vapor  pressure  of  solutions  explains  the  de- 
liquescence of  certain  substances.  Many  substances  absorb 
water  when  exposed  to  the  air,  become  moist,  and  sometimes 
even  dissolve  in  the  absorbed  water.  Calcium  chloride, 
potassium  carbonate,  zinc  chloride,  sodium  hydroxide, 
magnesium  chloride,  and  potassium  hydroxide  belong  to 
this  class.  This  property  is  called  deliquescence,  and  the 
substances  are  said  to  deliquesce,  or  to  be  deliquescent. 
Deliquescence  is  a  property  of  very  soluble  substances. 
Water  vapor  from  the  air  condenses  on  the  surface  and  pro- 
duces a  very  concentrated  solution,  which  has  a  vapor 
pressure  much  lower  than  the  average  pressure  of  the  water 


GENERAL  PROPERTIES  OF  WATER      75 

vapor  in  the  air.  The  solution,  therefore,  continues  to  take 
up  water  until  its  vapor  pressure  equals  the  partial  pressure 
of  the  water  vapor  in  the  air.  Common  salt,  or  sodium 
chloride,  often  deliquesces,  especially  in  damp  weather.  The 
deliquescence  is  due,  however,  to  the  presence  of  magnesium 
and  calcium  chlorides.  Sodium  nitrate  is  somewhat  deli- 
quescent, and  is  not  usually  used  in  the  manufacture  of 
gunpowder;  potassium  nitrate,  which  is  not  deliquescent,  is 
used  instead.  This  property  of  deliquescence  is  often  utilized 
in  the  laboratory  to  remove  water  vapor  from  gases,  calcium 
chloride  being  usually  employed  for  this  purpose.  Substances 
thus  used  are  often  called  drying  or  desiccating  agents. 

Thermal  Phenomena  of  Solution.  — The  process  of  solu- 
tion is  often  accompanied  by  an  appreciable  change  of  tem- 
perature. Thus,  when  sulphuric  acid  is  poured  into  water, 
heat  is  liberated.  With  relatively  large  quantities  of  acid 
the  heat  is  so  great  that  the  mixture  often  boils,  and  some- 
times the  hot  acid  is  spattered.  Hence,  the  acid  should  be 
added  slowly  to  the  water,  and  the  mixture  constantly  stirred. 
Other  substances  which  dissolve  with  the  liberation  of  heat 
are  fused  calcium  chloride,  potassium  hydroxide,  and  sodium 
hydroxide.  Some  substances  which  dissolve  with  a  fall  of 
temperature  (i.e.  with  absorption  of  heat)  are  crystallized 
calcium  chloride,  ammonium  nitrate,  ammonium  chloride, 
and  potassium  nitrate.  The  final  result  of  the  entire  act  of 
solution  is  doubtless  due  to  several  factors.  One  of  these  is 
the  change  in  volume  of  the  solute.  Thus,  a  very  soluble 
gas  becomes  greatly  condensed,  while  a  solid,  on  the  other 
hand,  occupies  a  much  larger  volume  after  solution  than 
before.  Another  important  factor  is  chemical  action  be- 
tween the  solute  and  solvent.  The  heat  liberated  or  ab- 
sorbed during  the  process  of  solution  is  called  heat  of  solutioa 
(See  Chapter  XL) 


76  INORGANIC   CHEMISTRY 

Solution  and  Chemical  Action.  —  When  a  substance  dis- 
solves, it  is  so  modified  that  it  can  participate  more  readily 
in  chemical  changes.  Hence,  solution  is  an  aid  to  chemical 
change.  Thus,  if  dry  tartaric  acid  and  sodium  bicarbonate 
are  mixed,  there  is  no  evidence  of  chemical  action;  but  when 
the  mixture  is  poured  into  water,  the  copious  evolution  of 
carbon  dioxide  gas  is  conclusive  evidence  of  a  chemical  change. 
Similarly,  when  a  dry  mixture  of  ferrous  sulphate  and  potas- 
sium ferrocyanide  is  poured  into  water,  the  immediate 
appearance  of  a  blue  precipitate  shows  that  water  was 
needed  for  the  chemical  change.  Solution  is  .such  an  im- 
portant aid  to  chemical  action  that  many  substances  em- 
ployed in  the  laboratory  are  in  solution,  and  many  processes 
in  chemistry  are  "wet"  processes.  (See  Chapter  IX.) 

The  Nature  of  Solution  has  long  been  a  subject  of  specula- 
tion and  study.  The  problem  as  a  whole  is  still  unsolved, 
though  much  light  has  been  thrown  upon  the  question  by 
recent  investigations.  (See  Chapter  IX.) 

PROBLEMS 

1.  If  1.5  gm.  of  crystallized  barium  chloride  loses  .22  gm.  when 
heated  to  constant  weight,  what  per  cent  of  water  of  crystallization 
does  it  contain  ? 

2.  If  2  gm.  of  another  lot  of  barium    chloride    loses  .295    gm., 
what  per  cent  of  it  was  water  of  crystallization? 

3.  A  tube  contains  97.2  cc.  of  gas  at  20.3°  C.  and  756  mm.,  and  the 
vapor  pressure  is  17.65.     What  is  the  volume  of  the  dry  gas  at  0°  C. 
and  760  mm.  ? 

4.  Reduce  the  following  as  in  Problem  3:    (a)  77  cc.,  17.5°  C., 
755   mm.,    14.89   a;    (6)    81.2   cc.,   746.8   mm.,    19.5°  C.,    16.87   a; 
(c)   100  cc.,  755.3  mm.,  18.5°  a,  15.85  a. 

5.  If  the  density  of  ice  is  .92,  what  volume  will  a  liter  of  water 
at  4°  C.  occupy  when  frozen?  Ans.  1.087  1. 

6.  How  much  water  (approximately)   is  contained  in   (a)   2  Ib. 


GENERAL  PROPERTIES  OF  WATER  77 

of  lobster,  (6)  56  Ib.  of  potatoes,  (c)  1  Ib.  of  tomatoes,  (d)  2  Ib.  of 
milk,  (e)  1  Ib.  of  white  bread,  (/)  a  human  body  weighing  150  Ib.  ? 

7.  If  a  dry  vessel  of  1000  cc.  capacity  has  a  drop  of  water  put 
into  it  at  20°  C.,  what  weight  of  water  will  evaporate?     (NOTE.  — 
One  liter  of  water  vapor  at  0°  C.  and  760  mm.  weighs  .8045  gin.) 

8.  50  cc.  of  dry  hydrogen  at  18.3°  C.  and  758.7  mm.  would 
occupy  what  volume  at  0°  C.  and  760  mm.  when  saturated  with 
water  vapor  ? 

9.  Plot  the  following  data  on  cross-section  paper  and  draw 
the  solubility  curve  of  the  substance :     Temperature  —  0,  10,  20, 
40,  55,  80 ;    corresponding  solubility  (i.e.  grams  soluble  in  100  gm. 
of  water)  —  .8,  .946,  1.18,  1.7,  2.1,  3.1. 

10.  As  in  Problem  9 :    Temperature  —  0,  10,  20,  30,  40,  50,  60, 
70,  80,  90,  100;    solubility  —  26.9,  31.5,  36.2,  40.9,  45.6,  50.3,  55, 
59.6,  64.2,  68.9,  73.8. 

11.  30  cc.  of  a  solution  weigh  33.315  gm.  and  give  8.865  gm.  of 
solid  on  evaporation.     Calculate  (a)  the  solubility  of  the  solid,  and 
(6)  the  weight  of  the  solid  in  100  cc.  of  the  solution. 

12.  If  a  cake  of  ice  weighs  280  kg.,  what  is  its  volume? 

13.  A  solution  has  a  specific  gravity  of  1.8.     How  many  cubic 
centimeters  of  water  must  be  added  to  a  liter  of  it  to  reduce  its 
specific  gravity  to  1.5? 

14.  By  use  of  the  solubility  curves  on  page  68  answer :   (a)  How 
many  grams  of  potassium  chloride  are  in  solution  at  10°,  20°,  25°, 
60°,   80°,   95°,    100°?     (6)  Compare   the   solubility   of   potassium 
chlorate,  potassium  bromide,  potassium  chloride,  and  potassium 
nitrate.     How  much  of  each  is  in  solution  at  30°,  50°,  90°?     (c)  As 
in   (6)  —  sodium  chloride,  ammonium  chloride,  lead  nitrate,  and 
sodium  nitrate. 


CHAPTER  VI 
Composition  of  Water  —  Hydrogen  Dioxide 

WATER  was  called  an  element  until  about  the  end  of  the 
eighteenth  century.  At  that  time  it  was  shown  to  be  a 
compound  of  hydrogen  and  oxygen.  Since  water  is  the  first 
chemical  compound  we  are  to  study,  special  attention  will 
be  paid  to  its  typical  characteristics. 

We  should  recall  at  this  point  the  essential  character- 
istic of  a  chemical  compound,  viz.  its  constituents  are 
elements  chemically  combined  in  a  fixed  ratio  by  weight. 

The  Composition  of  a  Compound  is  determined  either  by 
analysis  or  synthesis,  i.e.  by  taking  it  apart  or  putting  its 
parts  together.  Sometimes  both  methods  are  used,  since 
each  fortifies  the  other  and  strengthens  the  final  conclusion. 
These  methods  find  excellent  application  in  determining 
the  composition  of  water.  Analysis  and  synthesis  may  be 
qualitative  or  quantitative.  A  qualitative  experiment  is  a 
study  of  compounds  with  a  view  of  discovering  what  ele- 
ments or  groups  of  elements  they  contain.  A  quantitative 
experiment  is  an  accurate  determination  of  the  weight  or 
volume  of  the  constituents  of  a  compound.  Obviously,  a 
complete  study  of  the  composition  of  a  compound  requires 
both  kinds  of  tests,  the  qualitative  as  a  rule  preceding  the 
quantitative. 

Water  contains  Hydrogen.  —  When  steam  is  passed  over 
heated  metals,  hydrogen  is  liberated.  Lavoisier's  demon- 
stration of  this  fact  has  already  been  considered  (see  Prepara- 
tion of  Hydrogen).  The  fact  that  sodium  liberates  hydrogen 

78 


COMPOSITION   OF  WATER  79 

from  water  at  the  ordinary  temperature  has  also  been  dis- 
cussed (see  ibid.).  If  red  litmus  paper  is  put  into  the  water 
from  which  the  sodium  has  liberated  hydrogen,  the  litmus 
paper  becomes  blue.  This  change  of  color  from  red  to  blue 
shows  that  an  alkali  is  in  the  water,  because  alkalies  turn 
red  litmus  paper  blue.  The  alkali  is  sodium  hydroxide,  and 
it  may  be  obtained  as  a  white  solid  by  evaporating  the  water. 
Sodium  hydroxide  is  a  compound  of  sodium,  hydrogen,  and 
oxygen,  and  is  formed  by  replacing  half  of  the  hydrogen  of 
water  by  sodium.  Since  sodium  liberates  hydrogen  from 
water,  and  forms  at  the  same  time  a  compound  —  sodium 
hydroxide  —  containing  hydrogen,  the  hydrogen  in  water 
must  be  divisible  into  two  parts.  Now  if  .1  gm.  of  sodium 
is  allowed  to  act  upon  water,  48.22  cc.  of  hydrogen  are 
liberated;  and  if  the  sodium  hydroxide  thus  formed  is  dried 
and  heated  with  sodium,  48.22  cc.  more  of  hydrogen  are 
obtained.  This  shows  that  the  hydrogen  in  water  is  divisible 
into  two  equal  parts  —  a  fact  of  fundamental  importance. 

Water  contains  Oxygen.  —  The  fact  that  oxygen  is  a 
constituent  of  water  has  already  been  suggested ;  e.g.  (1)  by 
the  production  of  wateV  when  hydrogen  is  burned  in  air, 
(2)  by  the  formation  of  a  compound  of  iron  and  oxygen 
when  steam  is  passed  over  hot  iron,  and  (3)  by  the  formation 
of  sodium  hydroxide  when  sodium  interacts  with  water. 
These  proofs,  however,  are  indirect.  A  simple  direct  demon- 
stration of  the  presence  of  oxygen  in  water  can  be  made 
by  allowing  chlorine  water  to  stand  in  the  sunlight.  (Chlorine 
water  is  prepared  by  saturating  water  with  chlorine  gas  - 
an  element  to  be  studied  in  Chapter  XII.)  A  tube  about  a 
meter  long  is  completely  filled  with  chlorine  water,  the  open 
end  is  immersed  in  a  vessel  containing  the  same  solution 
and  the  whole  apparatus  is  placed  in  the  direct  sunlight. 
Bubbles  of  gas  soon  appear  in  the  liquid,  and  after  a  few 


80 


INORGANIC  CHEMISTRY 


hours  a  small  volume  of  gas  collects  at  the  top  of  the  tube. 
This  gas  can  be  shown  to  be  oxygen  by  the  usual  test,  viz. 
relighting  a  glowing  splinter  of  wood. 

Decomposition  of  Water.  —  When  steam  is  heated  to  a 
very  high  temperature  (about  2500°  C.),  it  decomposes  to 

a  very  slight  extent  into  hydro- 
gen and  oxygen.  This  method, 
however,  is  not  as  convenient  as 
the  one  in  which  an  electric  cur- 
rent is  used.  The  decomposi- 
tion of  water  by  electricity  is 
called  traditionally  the  elec- 
trolysis of  water,  though  we 
shall  see  later  (Chapters  IX  and 
XI)  that  the  process  is  more 
complex  than  the  apparent  dis- 
ruption of  water  into  hydrogen 
and  oxygen.  The  operation  is 
accomplished  in  a  special  form 
of  apparatus  devised  by  Hof- 
mann  and  shown  in  Figure  10. 
Pure  water  does  not  conduct 
electricity,  so  a  mixture  of  water 
(10  vols.)  and  sulphuric  acid 
(1  vol.)  is  poured  into  the  ap- 
paratus until  the  reservoir  is 
half  full  after  the  stopcocks  have  been  closed.  As  soon  as 
an  electric  battery  of  three  or  more  cells  is  connected  by 
wires  with  the  piece  of  platinum  near  the  bottom  of  each 
tube,  bubbles  of  gas  gather  on  the  platinum,  and  as 
the  action  proceeds,  the  bubbles  rise,  collect  in  the  upper 
part  of  the  tubes,  and  slowly  force  the  liquid  from  each 
tube  into  the  reservoir.  The  volume  of  gas  is  greater  in  one 


Fio.  10.  —  Special  form  of  Hof- 
mann  apparatus  for  the  elec- 
trolysis of  an  acid  solution  of 
water. 


COMPOSITION   OF   WATER 


81 


tube.  Assuming  that  the  tubes  have  the  same  diameter,  the 
volumes  are  in  the  same  ratio  as  their  heights,  which  will  be 
found  by  measurement  to  be  approximately  two  to  one. 
Appropriate  tests  show  that  the  gas  having  the  larger  volume 
is  hydrogen  and  the  one  having  the  smaller  volume  is  oxygen. 
Many  accurate  repetitions  of  this  experiment  have  shown 
that  only  hydrogen  and  oxygen  are  produced,  and  that  the 
volume  of  the  hydrogen  is  twice  that  of  the  oxygen. 

Water  was  first  decomposed  by  electricity  in  1800  by  Nicholson 
and  Carlisle.  Davy  confirmed  their  work  by  a  series  of  brilliant 
experiments  extending  through  a  period  of  six  years  (1800-1806). 

The  Quantitative  Composition  of  Water.  —  Decisive  evi- 
dence of  the  quantitative  composition  of  water  is  obtained 
by  a  determination  of  its 
volumetric  and  its  gravi- 
metric composition.  Volu- 
metric means  "  by  volume  " 
and  gravimetric  means 
"by  weight." 

The  Volumetric  Compo- 
sition of  Water  is  deter- 
mined by  exploding  a 
mixture  of  known  volumes 
of  hydrogen  and  oxygen 
in  a  eudiometer.  It  is  a 
method  of  synthesis. 

A  simple  sketch  of  a 
convenient  form  of  appa- 
ratus for  determining  the 

!  ,    .  ...  f  FIG.  11. — Apparatus  for  determining  the 

volumetric  composition  of        volumet^  cornposition  of  water. 
water  is  shown  in  Figure 

11.  The  essential  part  is  the  eudiometer,  F.  In  this 
glass  tube  the  gases  are  accurately  measured  and  exploded. 


82  INORGANIC  CHEMISTRY 

The  electric  spark  which  causes  the  explosion  is  obtained 
from  an  induction  coil  and  battery.  The  spark  leaps  across 
the  space  between  the  platinum  wires  at  the  top  of  the 
inside  of  the  eudiometer,  and  the  heat  produced  by  this 
spark  causes  the  hydrogen  and  oxygen  to  combine  and 
form  water.  Omitting  details,  oxygen  and  hydrogen  are 
introduced  separately  into  the  eudiometer  and  measured 
and  the  mixture  is  then  exploded;  after  the  explosion,  which 
is  indicated  by  a  slight  click  or  flash  of  light,  water  from  the 
reservoir,  E,  rushes  up  into  the  eudiometer.  The  water 
does  not  completely  fill  the  eudiometer,  because  an  excess 
of  one  gas  was  added.  This  additional  gas  takes  no  part  in 
the  chemical  change,  but  merely  serves  to  lessen  the  violence 
of  the  explosion,  which  otherwise  might  break  the  eudiom- 
eter. The  quantity  of  water  formed  by  the  union  of  the 
hydrogen  and  oxygen  is  too  minute  to  measure.  A  concrete 
illustration  will  make  the  process  more  intelligible.  Sup- 
pose the  volumes  of  the  participating  gases  after  reduction 
to  standard  temperature  and  pressure  and  the  dry  state  are 
as  follows  :  — 

Vol.  of  hydrogen  added 32.4  cc. 

Vol.  of  oxygen  added 12.3  cc. 

Vol.  of  residual  hydrogen 7.8  cc. 

The  actual  volume  of  hydrogen  which  combined  with  all 
the  oxygen  is  24.6  cc.  (i.e.  32.4  —  7.8).  Therefore  the  two 
gases  combined  in  the  ratio  of  24.6  to  12. 3;  or  2  to  1;  or,  as 
it  is  usually  stated,  two  volumes  of  hydrogen  combine  with 
one  volume  of  oxygen  to  form  water.  The  ratio  based  on 
the  most  painstaking  work  is  given  as  2.0027  to  1. 

The  discovery  of  the  volumetric  composition  of  water  was  not 
made  by  any  one  chemist.  Priestley,  about  1780,  noticed  that  when 
a  mixture  of  air  and  hydrogen  was  exploded,  "  the  inside  of  the  glass, 
though  clear  and  dry  before,  immediately  became  dewy."  Cavendish, 


COMPOSITION  OF  WATER 


83 


in  1781,  showed  that  when  a  mixture  of  two  parts  hydrogen  and  one 
part  oxygen  was  exploded,  nothing  but  water  was  formed.  Watt,  in 
1783,  was  the  first  to  state  that  water  is  a  compound,  though  he  per- 
formed no  experiments  and  probably  did  not  understand  the  real 
nature  of  its  constituents.  Lavoisier  in  the  same  year  verified  many 
facts  previously  noticed  but  not  completely  understood,  and  un- 
doubtedly first  clearly  recognized  and  stated  what  his  contempo- 
raries had  overlooked.  The  final  proof  of  the  volumetric  composi- 
tion of  water  was  an  accurate  verification  in  1805  by  Gay-Lussac 
and  Humboldt  of  the  previous  observation  that  two  volumes  of 
hydrogen  unite  with  one  volume  of  oxygen. 

The  Gravimetric  Composition  of  Water  is  determined  by 
passing  dry  hydrogen  over  copper  oxide.  The  method  de- 
pends upon  the  fact,  previously  stated,  that  many  oxides, 
such  as  those  of  lead,  copper,  and  iron,  when  heated  in  a 
current  of  hydrogen,  give  up  their  oxygen,  or,  chemically 
speaking,  these  oxides  are  reduced  to  metals.  By  this 
reduction  the  oxygen  of  the  oxide  combines  with  the  hydro- 
gen, thereby  forming  water,  which  is  collected  in  a  weighed 
tube,  while  the  metal  remains  behind  in  the  original  tube. 

A  sketch  of  the  apparatus  is  shown  in  Figure  12.  The  cop- 
per oxide  is  placed  in  the  combustion  tube,  CC,  which  is 


Fio.  12.  —  Apparatus  for  determining  the  gravimetric  composition  of  water. 

made  of  hard  glass.  The  Marchand  tube,  D,  which  is  filled 
with  calcium  chloride,  collects  and  retains  the  water  formed 
in  the  combustion  tube  by  the  reduction  of  the  copper 


84  INORGANIC  CHEMISTRY 

oxide.  The  tubes  A,  B,  and  E  keep  moisture  out  of  the 
apparatus.  The  experiment  is  easily  conducted.  Copper 
oxide  is  placed  in  the  combustion  tube,  which  is  then  care- 
fully weighed.  The  Marchand  tube,  filled  with  calcium 
chloride,  is  also  weighed.  After  the  other  tubes  are  properly 
filled  and  are  connected  as  shown  in  the  figure,  the  hydrogen 
generator  is  adjusted  so  that  a  slow  current  can  be  passed 
through  the  whole  apparatus  (from  left  to  right).  The 
combustion  tube  is  now  heated,  and  moisture  collects  in  it; 
as  the  heat  increases,  the  copper  oxide  glows,  and  the  mois- 
ture passes  into  the  Marchand  tube.  When  the  operation 
is  over  and  the  apparatus  is  cool  and  free  from  hydrogen,  the 
combustion  tube  and  Marchand  tube  are  weighed.  The  gain 
in  weight  of  the  Marchand  tube  is  the  weight  of  the  water 
formed,  while  the  loss  in  weight  of  the  combustion  tube  is 
the  weight  of  the  oxygen  removed  from  the  copper  oxide 
and  now  combined  with  the  hydrogen  in  the  form  of  water. 
An  example  will  make  this  clear.  Dumas,  who  first  did 
this  experiment  accurately,  found  substantially  that  the 
combustion  tube  lost  5.251  gm.  of  oxygen,  while  the  Mar- 
chand tube  absorbed  5.909  gm.  of  water.  Now  the  5.909 
gm.  of  water  contains  .658  gm.  of  hydrogen  (i.e.  5.909  - 
5.251).  But  .658  and  5.251  are  in  the  same  ratio  as  1  and 
7.98.  That  is,  water  contains  1  part  of  hydrogen  and  7.98 
parts  of  oxygen  by  weight.  This  ratio  is  very  nearly  1  to  8, 
and  the  gravimetric  composition  of  water  is  often  stated  as 
being  1  part  hydrogen  and  8  parts  oxygen.  Occasionally 
the  gravimetric  composition  is  stated  in  per  cent,  the  values 
being  11.18  per  cent  hydrogen  and  88.82  per  cent  oxygen. 
The  ratio  obtained  by  Dumas  so  long  ago  was  scarcely  ques- 
tioned until  recently.  We  now  know  from  an  exceptionally 
accurate  determination  by  Morley  that  the  ratio  is  1  to 
7.9395.  He  effected  a  complete  synthesis  of  water  in  which 
the  oxygen,  hydrogen,  and  water  were  weighed.  This  ratio 


COMPOSITION   OF  WATER 


85 


is  now  accepted  as  the  correct  one,  though  the  more  usual 
form  is  2  to  15.879.  The  apparatus  used  by  Morley  is  shown 
in  Figure  13.  It  was  first  weighed  vacuous  (i.e.  free  from 
air  or  other  gases).  The  tubes,  aa,  were 
then  connected  with  the  weighed  reservoirs 
of  oxygen  and  hydrogen,  and  the  oxygen 
was  introduced.  Sparks  were  next  passed 
between  the  platinum  wires,  cc,  and  the 
heat  ignited  the  hydrogen,  which  was  slowly 
admitted,  the  combination  of  the  gases 
taking  place  at  bb.  The  water  vapor 
condensed  in  the  tube  dd,  the  lower  por- 
tion of  which  was  immersed  in  water. 
The  combustion  of  the  hydrogen  was  con- 
tinued until  a  suitable  weight  of  water  was 
formed.  The  water  and  its  vapor  were 
then  converted  into  ice  by  putting  the 
apparatus  into  a  freezing  mixture;  the 
residual  mixture  of  gases  was  drawn  off 
and  analyzed,  passing  in  its  exit  through 
tubes  of  phosphorus  pentoxide  in  ee  which 
retained  all  traces  of  water.  The  whole 
apparatus  was  finally  weighed,  the  increase 
being  the  weight  of  the  water  formed  by  the  combination 
of  known  weights  of  hydrogen  and  oxygen. 

Summary.  —  Water  is  a  chemical  compound  of  hydrogen 
and  oxygen  combined  in  a  fixed  ratio  by  weight,  viz.  1  to 
7.9395;  they  are  also  combined  in  the  ratio  of  2.0027  to  1 
by  volume.  Usually  these  ratios  are  stated  approximately 
as  2  to  16  by  weight  and  2  to  1  by  volume. 

The  gravimetric  composition  of  water  was  first  determined  about 
1820  by  Berzelius  and  Dulong.  Their  work  was  verified  by  Dumas 
about  1842.  The  complete  synthesis  was  made  in  1895  by  Morley. 


FIG.  13.  —  Morley's 
apparatus  for  de- 
termining the  grav- 
imetric composi- 
tion of  water. 


86  INORGANIC  CHEMISTRY 


PROBLEMS  AND  EXERCISES 

1.  What  (a)  weight  and  (6)  volume  of  hydrogen  are  needed 
to  change  75  1.  of  oxygen  into  water?     (Assume  standard  condi- 
tions.) 

2.  What  (a)  weight  and  (6)  volume  of  oxygen  are  needed  to 
change  185  cc.  of  hydrogen  into  water  ?    (Assume  standard  conditions.) 

3.  Suppose   150  1.   of  water  are    decomposed    by  electricity. 
What  (a)  weight  and  (6)  volume  (at  0°  C.  and  760  mm.)  of  the 
products  are  formed? 

4.  How  many  (a)  grams  and   (6)  cubic  centimeters  (at  0°  C. 
and  760  mm.)  can  be  prepared  from  a  metric  ton  of  water? 

5.  Apply  Problem  4  to  oxygen.     What  would  be  the  volume 
of  oxygen  at  18°  C.  and  767  mm.  ? 

6.  If  20  gm.  of  water  were  produced  by  the  explosion  of  a  mix- 
ture of  hydrogen  and  oxygen,  what  volumes  of  the  dry  gases  were 
used  at  15°  C.  and  770  mm.  ? 

7.  How  many  grams  of  potassium  chlorate  are  needed  to  pro- 
duce enough  oxygen  for  the  complete  combustion  of  10  gm.  of 
hydrogen  ? 

8.  What  volume  of  (a)  hydrogen  and  (6)  oxygen,  both  at  12°  C. 
and  762  mm.,  can  be  obtained  by  decomposing  10  gm.  of  water 
by  electricity? 

9.  Morley  found  that  3.2645  gm.  of  hydrogen  combined  with 
25.9176  gm.  of  oxygen.     What  is  the  gravimetric  composition  of 
water  according  to  this  experiment  ? 

10.  How  would  you  prove  that  water  is  composed  of    only 
hydrogen  and  oxygen  ? 

11.  State  the  sources  of  error  in  determining  the  gravimetric 
composition  of  water  by  the  copper  oxide  method. 

12.  State  (a)  the  exact  gravimetric  and  (6)  the  exact  volumetric 
composition  of  water.     Describe  the  method  of  determining  each. 

HYDROGEN  DIOXIDE 

Hydrogen  Dioxide  is  a  liquid  composed  of  hydrogen  and 
oxygen.  But  the  proportion  of  the  constituents  is  not  the 
same  as  in  water.  It  contains  approximately  one  part  of 
hydrogen  and  sixteen  parts  of  oxygen  by  weight.  It  is 
often  called  hydrogen  peroxide,  because  its  relative  pro- 


COMPOSITION   OF  WATER  87 

portion  of  oxygen  is  greater  than  in   water  —  the  other 
hydrogen  oxide. 

Preparation.  —  It  is  prepared  by  treating  barium  dioxide 
(or  peroxide)  with  sulphuric  acid.  The  commercial  -solu- 
tion has  a  variable  strength,  and  usually  contains  3  per  cent 
or  more  of  hydrogen  dioxide. 

Properties.  —  Hydrogen  dioxide  has  a  sharp,  pungent 
odor  and  a  bitter,  metallic  taste.  The  concentrated  solu- 
tion is  a  syrupy  liquid,  but  the  commercial  solution  is  scarcely 
distinguishable  in  appearance  from  water.  It  is  an  unstable 
compound,  which  decomposes  slowly  at  the  ordinary  tem- 
perature, and  very  rapidly  if  heated.  The  dilute,  commercial 
solution  is  somewhat  stable,  but  heat  and  light  decompose 
it  completely  into  water  and  oxygen.  The  ease  with  which 
it  yields  oxygen  makes  it  a  good  oxidizing  agent.  Under 
certain  conditions  it  is  also  a  reducing  agent. 

Uses.  —  Dilute  solutions  are  used  extensively  to  bleach 
animal  and  vegetable  matter,  such  as  human  hair,  ostrich 
feathers,  fur,  silk,  wool,  cotton,  bone,  and  ivory.  It  is  also 
used  as  an  antiseptic  and  disinfectant  in  surgery  on  account 
of  its  oxidizing  properties.  Large  quantities  are  used  to 
restore  the  color  to  faded  paintings  —  a  use  suggested  by 
Thenard,  the  discoverer.  In  the  laboratory  it  is  a  service- 
able reagent. 


CHAPTER  VII 

Law,  Theory,  and  Hypothesis  —  Laws  of  Definite  and  Multiple 
Proportions  —  Atomic  Theory  —  Atoms  and  Molecules  — 
Symbols  and  Formulas  —  Equations 

Law  and  Theory.  —  We  discover  facts  by  observation  and 
experiment.  Phenomena  which  always  occur  under  the 
same  conditions  soon  become  well-established  facts.  Our 
knowledge  of  the  facts  which  have  some  relation  to  each 
other  is  often  summarized  in  a  brief  statement  called  a  law. 
A  law  is  not  only  an  epitome  of  the  uniform  behavior  of 
observed  facts.  It  is  also  a  statement  which  permits  us  to 
predict  occurrences  under  like  conditions;  for  if  a  law  is  valid 
for  many  observed  cases,  we  conclude  that  it  will  cover  future 
cases.  The  essential  feature  of  a  law  is  universal  validity. 

The  ultimate  cause  of  scientific  facts  is  unknown.  The 
explanation  we  give  of  facts,  especially  of  groups  of  related 
facts,  is  called  a  theory.  Laws  are  statements  about  facts; 
theories  are  statements  of  the  supposed  cause  of  facts. 
Laws  seldom  change,  but  theories  are  often  modified.  Laws 
are  the  result  of  experiment,  but  theories  are  the  outcome  of 
mental  operations.  We  accept  a  certain  theory  until  a  more 
satisfactory  one  is  proposed.  Sometimes  experiment  yields 
results  which  need  further  examination  or  are  beyond  the 
realm  of  our  present  experimental  skill.  The  temporary 
explanation  or  supposition  we  make  in  such  a  case  as  a  guide 
in  further  investigation  is  called  an  hypothesis.  An  hypothe- 
sis is  often  the  forerunner  of  a  theory,  and  they  are  riot  always 
sharply  differentiated. 

Laws,  theories,  and  hypotheses  are  of  great  service  in 

88 


LAW,   THEORY,   AND   HYPOTHESIS  89 

chemistry,  since  they  permit  us  to  gather  into  intelligible 
statements  our  knowledge  of  a  vast  number  of  related  facts, 
as  well  as  assist  us  in  discovering  new  facts  and  interpreting 
the  phenomena  of  nature.  In  this  chapter  we  shall  discuss 
two  laws  and  one  theory  that  are  of  great  importance  in 
chemistry. 

Law  of  Definite  Proportions  by  Weight.  —  When  the  metal 
magnesium  is  heated  in  the  air,  it  burns  with  a  dazzling  flame 
and  yields  a  grayish  powder,  due  to  combination  with  oxy- 
gen. If  magnesium  is  heated  so  that  the  product  cannot 
escape,  a  remarkable  relation  between  their  weights  is  re- 
vealed. In  order  to  burn  completely  1.52  gm.  of  magnesium, 
1  gm.  of  oxygen  is  necessary;  and  the  product,  magnesium 
oxide,  weighs  2.52  gm.  This  product  contains,  therefore, 
60.317  per  cent  of  magnesium  and  39.682  per  cent  of  oxygen. 
Accurate  repetitions  of  this  experiment  have  shown  that  this 
proportion  by  weight  is  fixed  and  definite.  Again,  if  a  weighed 
quantity  of  potassium  chlorate  is  decomposed,  31.903  per 
cent  of  potassium,  28.932  per  cent  of  chlorine,  and  39.164 
per  cent  of  oxygen  are  always  obtained.  This  means  that 
the  proportion  of  potassium,  chlorine,  and  oxygen  which 
makes  up  potassium  chlorate  is  fixed  and  definite.  Experi- 
ments similar  to  these  show  that  in  all  of  the  chemical 
compounds  which  have  been  examined  the  different  con- 
stituents are  always  present  in  a  definite  and  unvarying 
proportion  by  weight.  There  are  no  exceptions  to  this 
general  fact.  This  constancy  of  proportion  in  chemical  com- 
pounds is  stated  as  the  Law  of  Definite  Proportions  by 
Weight,  thus:  — 

A  given  chemical  compound  always  contains  the  same  ele- 
ments in  the  same  proportion  by  weight. 

This  law  is  one  of  the  fundamental  laws  of  chemistry.  It 
is  so  confidently  believed  that  if  the  composition  of  a  com- 


90  INORGANIC  CHEMISTRY 

pound  is  found  by  analysis  to  vary,  chemists  conclude  that 
the  experimental  work  is  incorrect  or  that  the  compound  is 
impure.  The  law  was  established  as  the  outcome  of  a  con- 
troversy between  two  French  chemists,  Proust  (1755-1826) 
and  Berthollet  (1748-1822).  The  discussion  lasted  from 
1799  to  1806.  Berthollet  believed  that  compounds  might 
have  a  varying  composition.  Indeed,  by  his  experiments 
he  detected  " gradual  changes"  in  composition.  But  Proust 
showed  that  Berthollet  analyzed  mixtures  and  not  com- 
pounds. (In  a  mixture  the  components  may  be  present 
in  any  proportion.)  Subsequent  experiments  have  only 
strengthened  our  confidence  in  this  law. 

Law  of  Multiple  Proportions.  —  Proust  showed  that  some 
elements  combine  in  more  than  one  proportion  by  weight, 
and  thereby  produce  two  or  more  distinct  compounds.  But 
he  failed  to  notice  that  if  the  weight  of  one  element  is  adopted 
as  constant,  the  varying  weights  of  the  other  element  or  ele- 
ments are  in  a  simple  multiple  relation  to  each  other.  Dalton 
discovered  this  general  fact  about  1804.  The  composition 
of  compounds  is  usually  expressed  in  per  cent;  but  such 
expressions  in  a  series  of  compounds  reveal  nothing  about 
multiple  relations.  If,  however,  a  constant  weight  of  one 
constituent  is  adopted  as  a  basis  of  comparison,  and  the  com- 
position of  the  series  of  compounds  is  expressed  in  terms  of 
this  weight,  then  the  simple  multiple  relation  which  exists 
between  the  weights  of  the  other  constituent  or  constituents 
is  clearly  seen.  For  example,  no  multiple  relation  is  evident 
from  the  statement  that  two  compounds  contain  respec- 
tively 27.27  and  42.8571  per  cent  of  carbon  and  72.72  and 
57.1428  per  cent  of  oxygen.  But  if  in  expressing  the  com- 
position of  these  compounds  we  adopt  some  convenient 
number,  such  as  1  for  the  weight  of  carbon  in  each  com- 
pound, the  weights  of  oxygen  will  be  in  the  simple  integral 


LAW,   THEORY,   AND   HYPOTHESIS 


91 


ratio  of  2  to  1.  The  five  compounds  of  oxygen  and  nitrogen, 
which  will  soon  be  studied,  aptly  illustrate  this  fact  of  mul- 
tiple proportions :  — 

TABLE  TO  ILLUSTRATE  MULTIPLE  PROPORTIONS 


COMPOSITION  IN 

ADOPTED 

PER  CENT 

WEIGHT 

NAME 

Nitrogen 

Oxygen 

Nitrogen 

Nitrogen 

Oxygen 

Nitrous  Oxide  .     .     . 

63.636 

36.363 

1 

.57 

Nitric  Oxide    .     .     . 

46.666 

53.333 

1 

1.14 

Nitrogen  Trioxide 

36.842 

63.157 

1 

1.71 

Nitrogen  Peroxide     . 

30.434 

69.565 

1 

2.28 

Nitrogen  Pentoxide  . 

25.925 

74.074 

1 

2.85 

From  this  table  it  is  clear  that  the  weights  of  oxygen  com- 
bined with  the  same  weight  of  nitrogen  are  as  1  :  2  :  3  :  4  :  5 ; 
i.e.  they  are  in  a  relation  to  each  other  which  can  be  expressed 
by  small  integral  numbers.  The  same  simple  ratio  would  be 
obtained  if  any  other  value  were  substituted  for  1,  and  similar 
tables  may  be  worked  out  with  other  series  of  compounds. 

The  general  fact  of  multiple  proportions  is  stated  as  the 
Law  of  Multiple  Proportions,  thus:  — 

When  two  or  more  elements,  a,  b;  c,  etc.,  unite  to  form  a  series 
of  compounds,  a  fixed  weight  of  a  always  combines  with  such 
weights  of  b  (as  well  as  of  c,  etc.),  that  the  ratio  between  these 
different  weights  can  be  expressed  by  small  (usually),  whole 
numbers. 

This  law,  like  the  law  of  definite  proportions,  is  a  funda- 
mental law.  And  together  with  the  law  of  the  conservation 
of  matter  they  have  profoundly  influenced  the  theoretical 
and  practical  progress  of  chemistry. 

The  Atomic  Theory.  — The  laws  just  discussed  state  in 
condensed  form  certain  general  facts  about  the  quantitative 


92  INORGANIC   CHEMISTRY 

aspects  of  chemical  change.  They  point  to  the  existence  of 
chemical  units  which  participate  in  chemical  changes  with' 
out  alteration  of  weight.  But  we  have  no  means  of  detect- 
ing or  separating  these  chemical  units.  Nevertheless,  in  order 
to  provide  a  mental  picture  of  this  quantitative  feature  of 
chemical  change  a  theory  has  been  proposed.  It  is  called 
the  atomic  theory  and  was  announced  in  approximately  its 
present  form  by  Dalton,  an  English  chemist,  about  1805. 
According  to  this  theory,  (1)  chemical  elements  and  com- 
pounds consist  ultimately  of  a  vast  number  of  very  small 
particles  called  atoms;  (2)  chemical  change  is  union,  separa- 
tion, or  exchange  of  undivided  atoms;  (3)  atoms  of  the  same 
chemical  element  are  alike  and  have  an  unvarying  weight 
called  the  atomic  weight;  (4)  atoms  of  different  elements 
differ  from  each  other  in  weight.  The  atomic  theory  means 
in  a  few  words  that  matter  consists  of  atoms  which  are  en- 
dowed with  a  weight  characteristic  of  each  element  and  which 
remain  undivided  in  chemical  changes. 

This  theory,  it  will  be  observed,  deals  with  the  nature  of 
matter  and  with  the  quantitative  aspects  of  chemical  change. 
It  does  not  state  facts  nor  does  it  make  facts  more  valid. 
It  serves  merely  to  assign  some  explanation,  more  or  less 
detailed,  to  facts  and  laws  already  formulated.  Let  us 
apply  the  atomic  theory  to  certain  facts,  i.e.  restate  these 
facts  in  terms  of  the  theory.  Before  proceeding  with  this 
application,  however,  it  will  be  necessary  to  make  a  pre- 
liminary distinction  between  an  atom  and  a  molecule.  About 
the  time  of  Dalton  (1766-1844)  the  term  particle  was  used  to 
include  both  atom  and  molecule.  But  they  are  not  iden- 
tical. An  atom  is  the  smallest  particle  of  an  element  that 
participates  in  a  chemical  change.  Molecules  are  particles 
which  consist  of  two  or  more  atoms  chemically  combined ; 
atoms  are  alike  in  molecules  of  elements  but  different  in  mole- 
cules of  compounds.  Thus,  the  smallest  particle  of  copper 


LAW,  THEORY,  AND  HYPOTHESIS  93 

which  participates  in  chemical  changes  is  an  atom,  and  the 
smallest  particle  of  water  is  a  molecule,  which  consists  of 
hydrogen  and  oxygen  atoms  chemically  combined.  A  fuller 
discussion  of  atoms  will  soon  be  given.  But  this  preliminary 
distinction  will  permit  us  to  resume  the  application  of  the 
atomic  theory  to  certain  facts.  An  appreciable  mass  of  the 
element  copper  consists  of  many  millions  of  atoms  of  copper, 
all  alike  —  all  having  the  same  unvarying  weight.  Sodium, 
oxygen,  hydrogen,  sulphur,  carbon,  and  all  the  other  elements 
likewise  consist  of  atoms,  but  the  weight  of  the  sodium  atom 
differs  from  the  weight  of  the  atom  of  copper,  carbon,  sulphur, 
and  all  the  other  elements.  Again,  when  a  chemical  change 
occurs  between  copper  and  sulphur,  for  example,  atoms  of 
copper  combine  with  atoms  of  sulphur  and  produce  molecules 
of  a  compound  called  copper  sulphide.  And  this  combining 
of  atoms  into  molecules  continues  until  all  the  copper  atoms 
or  the  sulphur  atoms,  or  under  special  conditions  the  atoms 
of  both  substances,  have  been  used.  Furthermore,  this 
chemical  change  takes  place  not  only  between  vast  numbers 
of  atoms,  but  the  quantitative  aspects  of  this  multitude  of 
changes  conform  to  the  atomic  theory.  This  latter  point 
needs  explanation,  because  it  emphasizes  the  chief  feature 
of  the  atomic  theory,  viz.  agreement  with  certain  fundamental 
laws  of  chemical  change.  These  laws  are  the  law  of  the  con- 
servation of  matter,  the  law  of  definite  proportions,  and  the 
law  of  multiple  proportions.  (1)  According  to  the  atomic 
theory  the  weights  of  atoms  do  not  change ;  all  other  prop- 
erties may  be  temporarily  lost  or  buried,  but  the  weight  is  re- 
tained throughout  all  chemical  changes  however  complex.  It 
is  obvious  that  if  atoms  never  change  their  weights,  the  total 
weight  of  matter  in  a  chemical  change  is  unvaried,  and  the  law 
of  the  conservation  of  matter  follows  as  a  natural  consequence. 
(2)  Again,  according  to  the  atomic  theory,  when  magnesium 
combines  with  oxygen,  molecules  of  magnesium  oxide  are 


94  INORGANIC  CHEMISTRY 

formed  by  the  union  of  some  whole  number  of  atoms  of  mag- 
nesium with  some  whole  number  of  atoms  of  oxygen.  Each 
molecule  of  magnesium  oxide  would  therefore  consist  of  one 
or  more  atoms  of  magnesium  united  with  one  or  more  atoms 
of  oxygen,  and  the  composition  of  each  molecule  of  mag- 
nesium oxide  would  be  definite;  i.e.  each  molecule  would 
contain  the  same  elements  united  in  a  definite  proportion  by 
weight.  In  other  words,  magnesium  oxide  would  always  be 
found  to  consist  of  a  certain  per  cent  of  magnesium  and  a 
certain  per  cent  of  oxygen.  Hence  the  atomic  theory  har- 
monizes with  the  law  of  definite  proportions.  (3)  Finally,  the 
atomic  theory  conforms  to  the  law  of  multiple  proportions. 
The  number  of  atoms  of  the  combining  elements  may  be 
the  same  or  different.  That  is,  the  ratio  of  combination 
may  be  1  to  1,  1  to  2,  2  to  3,  etc.  According  to  the  atomic 
theory  atoms  are  transferred  as  wholes;  i.e.  in  chemical 
changes  there  are  no  fractions  of  atoms.  Therefore,  the 
proportions  of  the  weights  of  different  elements  in  a  series  of 
compounds  must  be  simple  proportions;  all  multiple  rela- 
tions will  be  expressible  by  whole  numbers.  An  illustration 
will  make  this  point  clear.  There  are  two  compounds  of 
carbon  and  oxygen.  Analysis  shows  that  in  the  one  contain- 
ing the  less  oxygen  the  ratio  of  the  weight  of  carbon  to  oxy- 
gen is  3  to  4,  and  in  the  other  3  to  8.  The  second  compound 
contains  twice  as  much  oxygen  as  the  first;  i.e.  its  molecule 
contains  twice  as  many  atoms  of  oxygen  as  a  molecule  of  the 
first.  (Subsequently  it  will  be  shown  that  the  first  com- 
pound is  carbon  monoxide  and  contains  one  atom  of  oxygen 
in  each  molecule,  while  the  second  is  carbon  dioxide  and 
contains  two  atoms  of  oxygen  in  each  molecule.)  In  other 
words,  the  weights  of  oxygen  in  this  series  of  compounds  are 
in  the  simple  ratio  of  1  to  2.  The  same  line  of  reasoning  can 
be  applied  to  other  series  of  compounds  whatever  the  ratio 
of  combination.  Hence  the  atomic  theory  harmonizes  with 
the  law  of  multiple  proportions. 


LAW,  THEORY,   AND   HYPOTHESIS  95 

Atoms  and  Molecules.  —  It  should  not  be  overlooked  that 
the  laws  of  definite  and  multiple  proportions  deal  with  facts, 
while  the  atomic  theory  deals  with  conceptions  which  may 
be  true,  but  which  we  cannot  prove  to  be  true.  We  often 
speak  of  atoms  as  if  they  could  be  perceived  by  the  senses, 
but  we  do  so  simply  because  such  methods  of  expression  help 
us  describe,  study,  and  interpret  chemical  changes.  We 
also  describe  elements  as  if  they  consisted  of  free  atoms,  but 
atoms  do  not,  as  a  rule,  exist  in  the  uncombined  state.  As 
soon  as  atoms  are  freed  from  combination,  they  at  once  unite 
with  some  other  atom  or  atoms.  The  particles  which  make 
up  oxygen  gas  are  not  atoms,  but  a  group  or  chemical  com- 
bination of  atoms.  Groups  of  chemically  combined  atoms 
are  called  molecules.  The  molecule  of  an  element  consists 
of  atoms  of  one  kind  only,  but  the  molecule  of  a  compound 
consists  of  atoms  of  two  or  more  different  kinds.  For 
example,  a  molecule  of  oxygen  consists  of  atoms  of  the  ele- 
ment oxygen,  while  a  molecule  of  the  compound  water  con- 
sists of  atoms  of  oxygen  and  hydrogen.  (A  molecule  of  a  few 
elements  contains  only  one  atom.)  Atoms  are  the  chemical 
constituents  of  molecules.  They  are  the  smallest  particles 
of  the  elements  which  are  known  to  participate  in  chemical 
changes. 

Our  views  regarding  molecules  are  based  on  extensive  in- 
vestigation of  the  properties  of  gases.  The  molecule  is  often 
spoken  of  as  the  physical  unit,  because  in  most  physical 
changes  molecules  are  not  decomposed.  W7hereas  the  atom 
is  called  the  chemical  unit  because  it  is  the  part  of  a  molecule 
which  as  a  rule  is  transferred  unchanged  in  chemical  changes. 
Molecules  will  be  discussed  again.  (See  Chapter  XIV.) 

Although  the  atom  is  conceived  to  pass  as  a  whole  from 
compound  to  compound,  it  should  not  be  inferred  that 
atoms  cannot  be  decomposed  under  any  conditions.  The 
phenomena  exhibited  by  compounds  of  radium  show  that 


96  INORGANIC  CHEMISTRY 

there  are  particles  smaller  than  atoms.  (See  Radio-activity.) 
These  very  small  particles  are  called  corpuscles  or  electrons. 
However,  the  atom  is  the  chemical  unit,  and  whether  a 
single  individual  or  a  group  of  smaller  individuals,  its  weight 
is  not  altered  in  chemical  changes. 

Atomic  Weights.  —  The  essential  property  of  matter  is 
weight.  According  to  the  atomic  theory,  different  kinds  of 
atoms  have  different  weights.  But  the  absolute  weight 
of  an  atom  cannot  be  directly  determined  by  any  instru- 
ments available.  We  can,  however,  find  the  relative  weight 
of  an  atom ;  that  is,  how  many  times  heavier  one  atom  is  than 
another  atom.  These  relative  weights  are  called  the  atomic 
weights  of  the  elements.  If  we  should  adopt  1  as  the  weight 
of  an  atom  of  hydrogen,  the  weights  of  atoms  of  other  ele- 
ments could  be  readily  expressed  in  terms  of  this  standard. 
Thus,  the  atomic  weight  of  sodium  would  be  22.88,  of  oxy- 
gen 15.88,  of  carbon  11.9,  etc.;  that  is,  an  atom  of  sodium 
weighs  twenty-three  times  as  much  as  an  atom  of  hydrogen, 
etc.  For  many  years  hydrogen  (=1)  was  the  standard. 
But  for  scientific  reasons  oxygen  is  now  the  standard,  and 
16  is  adopted  as  its  atomic  weight  instead  of  15.88.  This 
change  does  not  alter  any  facts  ;  it  merely  changes  slightly 
the  numerical  values  of  the  atomic  weights.  Their  relation 
to  each  other  is  not  changed.  The  atomic  weight  of  hydro- 
gen becomes  1.008,  if  oxygen  equals  16,  and  others  are  pro- 
portionally changed.  The  atomic  weights  are  real  weights 
because  they  are  found  by  experiment.  It  should  be  con- 
stantly borne  in  mind,  however,  that  they  are  relative 
weights.  That  is,  the  atomic  weight  of  sodium,  for  example, 
is  23.00,  not  23.00  gm.  or  any  other  absolute  weight,  but 
23.00,  if  the  atomic  weight  of  oxygen  is  16. 

The  exact  determination  of  the  atomic  weight  of  an  ele- 
ment is  a  difficult  operation.  Many  principles  influence  the 


LAW,   THEORY,    AND   HYPOTHESIS  97 

final  selection  of  the  number  to  be  adopted  as  the  atomic 
weight.  This  subject  is  discussed  in  Chapter  XIV.  A 
complete  table  of  the  atomic  weights  is  given  in  the  Appendix, 
§  5.  Exact  and  approximate  values  of  the  atomic  weights 
of  the  important  elements  can  be  found  in  a  table  on  the 
inside  of  the  back  cover,  and  for  most  purposes  these  approx- 
imate values  can  be  used,  e.g.  in  solving  the  problems  in  this 
and  subsequent  chapters. 

Chemical  Symbols,  which  were  mentioned  in  Chapter  I, 
represent  single  atoms.  Thus,  H  represents  one  atom  of 
hydrogen,  O  one  atom  of  oxygen,  N  one  atom  of  nitrogen. 
If  more  than  one  uncombined  atom  is  to  be  designated,  the 
proper  numeral  is  placed  before  the  symbol,  thus :  — 

2  H  means  2  uncombined  atoms  of  hydrogen, 

3  O  means  3  uncombined  atoms  of  oxygen, 

4  P  means  4  uncombined  atoms  of  phosphorus. 

But  if  we  wish  to  represent  the  atoms  as  in  chemical  com- 
bination, either  with  themselves  or  with  other  atoms,  then  a 
subscript  is  used  instead  of  a  coefficient,  thus :  — 

H2  means  2  atoms  of  hydrogen  in  combination, 
N3  means  3  atoms  of  nitrogen  in  combination, 
P4  means  4  atoms  of  phosphorus  in  combination. 

Symbols  not  only  represent  atoms,  but  they  also  express 
atomic  weights.  Thus,  0  represents  one  atom  of  oxygen,  but 
it  also  means  that  this  atom  weighs  16.  Similarly,  K  rep- 
resents one  atom  of  potassium,  which  weighs  39.10. 

Chemical  Formulas. — A  formula  is  a  group  of  symbols 
which  is  designed  to  express  the  composition  of  a  compound. 
A  given  compound,  as  we  have  already  seen,  has  a  definite 
composition.  In  other  words,  a  molecule  of  a  given  com- 


98  INORGANIC  CHEMISTRY 

pound  always  contains  the  same  number  and  the  same  kind 
of  atoms ;  and  since  the  molecules  are  alike,  the  composition 
of  one  molecule  is  the  same  as  the  composition  of  the  com- 
pound. In  writing  a  formula,  the  symbols  of  the  different 
atoms  making  up  a  molecule  of  the  compound  are  placed 
side  by  side.  Thus  H20  is  the  formula  of  water,  because  one 
molecule  of  this*  compound  consists  of  2  atoms  of  hydrogen 
and  1  atom  of  oxygen.  Similarly,  KC103  is  the  formula  of 
potassium  chlorate.  These  symbols  might  be  written  in  a 
different  order,  but  usage  has  determined  the  order  in  most 
cases.  A  formula,  as  just  stated,  represents  one  molecule. 
Hence,  KC1O3  represents  one  molecule  of  potassium  chlorate, 
and  means  that  the  molecule  of  this  compound  contains  1 
atom  each  of  potassium  and  chlorine  and  3  atoms  of  oxygen. 
If  we  wish  to  designate  several  molecules,  the  proper  nu- 
meral is  placed  before  the  formula,  thus  :  — 

2  KC103  means  2  molecules  of  potassium  chlorate, 

3  H2O  means  3  molecules  of  water, 

4  H2S04  means  4  molecules  of  sulphuric  acid. 

In  certain  compounds  some  of  the  atoms  act  like  a  single 
atom  in  chemical  changes.  This  fact  is  often  expressed  by 
inclosing  the  group  of  atoms  in  a  parenthesis,  or  by  separat- 
ing it  from  the  rest  of  the  formula  by  a  period.  Thus,  the 
formula  of  ammonium  nitrate  is  (NH4)NO3,  and  the  formula 
of  alcohol  is  C2H5 .  OH.  Water  of  crystallization  is  usually 
indicated  in  the  same  manner,  CuSO4.5  H20  being  the  for- 
mula of  crystallized  copper  sulphate.  The  period  is  some- 
times omitted,  especially  if  the  composition  of  the  compound 
is  well  understood.  If  a  group  of  atoms  is  to  be  multiplied, 
it  is  placed  within  a  parenthesis.  Thus,  the  formula  of  lead 
nitrate  is  Pb(N03)2.  This  means  that  the  group  NO3  is  to 
be  multiplied  by  2.  That  is,  lead  nitrate  might  be  expressed 
by  PbN3O6,  but  for  reasons  which  will  be  given  later  the 


LAW,   THEORY,   AND   HYPOTHESIS  99 

formula  Pb(N03)2  is  used.     The  expression  2  Pb(NO3)2  means 
that  the  formula  Fb(NO3)2  is  to  be  multiplied  by  2. 

Symbols  and  formulas  are  sometimes  used  to  represent  an  in- 
definite amount  of  an  element  or  compound.  Thus,  O  is  often  used 
to  designate  oxygen  and  H2SO4  sulphuric  acid.  They  are  often  thus 
used  to  label  bottles  in  a  laboratory.  Such  a  departure  from  accuracy 
should  not  be  allowed  to  obscure  their  real  meaning. 

The  complete  significance  of  symbols  and  formulas  can  be  grasped 
only  by  their  intelligent  use.  They  should  not  be  committed  to 
memory  slavishly.  It  is  desirable,  however,  to  learn  the  common 
ones  while  the  substances  they  represent  are  being  studied,  and 
leave  the  consideration  of  their  relations  until  the  needed  facts  have 
accumulated. 

Molecular  Weights.  — Since  atoms  combine  to  form  mole- 
cules, a  molecular  weight  is  the  sum  of  the  weights  of  the 
atoms  in  the  molecule.  A  molecule  of  nitric  acid  contains 

1  atom  each  of  hydrogen  and  nitrogen,  and  3  atoms  of  oxy- 
gen; hence  its  molecular  weight  is  1 +  14 .01 +  (16x3)  =63. 01. 
Given  the  formula,  the  molecular  weight  is  easily  found  by 
adding  the  atomic  weights.     Just  as  a  symbol  stands  for  an 
atomic  weight,  so  a  formula  expresses  a  molecular  weight. 
Molecular  weights  are  real  numbers,  but  they  are  not  found 
experimentally  by  merely  adding  the  atomic  weights.     For 
convenience  we  may  add  the  atomic  weights,  but  historically 
and  experimentally  molecular  weights  preceded  exact  atomic 
weights.     Many  facts  and  principles  determine  the  selection 
of  the  molecular  weight  of  a  compound.     These  are  discussed 
in  Chapter  XIV.     For  the  present,  empirical  knowledge  is 
sufficient.     That  is,  it  will  answer  all  purposes  to  find  the 
molecular   weight   of   a   compound   by   adding   the   atomic 
weights  corresponding  to  the  number  of  atoms  in  the  formula 
of  a  molecule  of  the  given  compound.     By  this  somewhat 
arbitrary  procedure  the  molecular  weight  of  water  (H2O)  is 

2  +  16  =  18;  the  weight  of  2  H2O  is  2(2  +  16)  =  36.     Simi- 
larly, the  molecular  weight  of  lead  nitrate  (Pb(NO8)2)  is  207.20 


100  INORGANIC  CHEMISTRY 

4-  2(14.01  +  48)  =  331.22;  the  weight  of  two  molecules  of 
lead  nitrate  (2  Pb(N03)2)  is  2  X  331.22  =  662.44.  Also,  the 
molecular  weight  of  sulphuric  acid  (H2SO4)  is  (2X1)+  32.06 

+  (4  X  16)  =  98.06.  Other  molecular  weights  may  be  calcu- 
lated in  the  same  way. 

It  should  be  noted  that  the  molecular  weight  of  a  com- 
pound, like  the  atomic  weight  of  an  element,  is  a  relative 
weight.  That  is,  the  molecular  weight  of  water  is  not  18 
gm.,  but  18,  if  the  atomic  weight  of  oxygen  is  16. 

Calculations  based  on  Formulas  and  Molecular  Weights.  — • 
It  is  evident  from  the  foregoing  paragraphs  that  there  is  a 
rigid  connection  between  the  molecular  weight  and  the 
formula  of  a  compound.  Since  the  formula  expresses  the 
composition  of  a  compound  by  means  of  small  integral  num- 
bers representing  the  ratio  of  the  atomic  weights  in  a  mole- 
cule, it  is  possible  to  calculate  (1)  the  composition  in  per  cent, 
if  the  formula  is  known,  and  (2)  the  formula,  if  the  compo- 
sition in  per  cent  is  known.  Composition  in  per  cent,  or,  as 
it  is  usually  designated,  percentage  composition,  is  readily 
calculated  from  the  formula  of  a  compound.  Let  us  take 
a  concrete  case.  The  formula  of  sulphuric  acid  is  HaS04. 
This  formula  represents  a  molecular  weight  of  98,  i.e. 
2  4-  32  -f  64  =  98  (using  approximate  atomic  -weights).  Now 
if  the  respective  parts  of  hydrogen,  sulphur,  and  oxygen  (viz. 
2,  32,  64)  are  divided  by  98  and  the  quotient  then  multiplied 
by  100  (e.g.  •£%  X  100),  the  product  is  the  per  cent  of  each 
element  in  sulphuric  acid.  It  is  sometimes  more  convenient 
to  solve  the  problem  by  a  proportion.  Thus,  the  proportions 
for  the  percentage  composition  of  sulphuric  acid  are :  — 

2  :  98  :  :  X  :  100;  X  =    2.04  per  cent  of  hydrogen. 
32  :  98  :  :  X  :  100;  X  =  32.65  per  cent  of  sulphur. 
64  :  98  :  :  X  :  100;  X  =  65.31  per  cent  of  oxygen. 
Total,  100.00  per  cent. 


LAW,   THEORY,   AND   HYPC'FHSisi&  *01 

By  the  same  method"  the  percentage  composition  of  any  com- 
pound can  be  calculated.  The  calculation  of  a  formula, 
when  the  atomic  weights  and  the  percentage  composition 
are  known,  is  practically  the  converse  of  the  above  process; 
i.e.  it  is  simply  the  process  of  finding  the  small  integral 
numbers  which  are  in  the  same  ratio  as  the  numbers  express- 
ing the  composition.  Suppose  we  know  the  composition 
of  sulphuric  acid  to  be  hydrogen  =  2.04  per  cent,  sulphur 
32.65  per  cent,  oxygen  65.31  per  cent.  If  the  percentage  of 
each  element  is  divided  by  the  corresponding  atomic  weight, 
the  quotients  are  2.04,  1.02,  and  4.08.  Reducing  these  quo- 
tients to  integral  numbers  (by  dividing  by  1.02),  the  final 
quotients  are  2, 1,  4.  But  these  quotients  represent  the  ratio 
of  the  atomic  weights  in  a  molecule;  that  is,  the  relative 
number  of  atoms  of  each  element  in  a  molecule.  Therefore 
the  formula  of  sulphuric  acid  must  be  H-jSO^  The  formula 
of  a  compound  calculated  by  this  method  is  its  simplest 
formula.  (See  also  Determination  of  Formulas  of  Com- 
pounds, Chapter  XIV.) 

Chemical  Equations.  —  When  substances  interact  chem- 
ically, the .  definite  chemical  transformation  is  called  a 
reaction.  We  have  already  seen  that  reactions  can  be  ex- 
pressed by  equations.  Thus,  in  Chapter  II  two  of  the  reac- 
tions involved  in  the  preparation  of  oxygen  were  expressed 
as  follows:  — 

Barium  Oxide  +  Oxygen  =  Barium  Dioxide. 

Barium  Dioxide  =  Barium  Oxide  +  Oxygen. 

It  was  stated  at  that  point  that  these  equations  are  crude 
forms  of  chemical  equations.  We  can  remodel  these  prelim- 
inary equations  by  using  symbols  and  formulas  in  place  of 
words,  the  equations  just  given  then  becoming  :  - 

BaO  +  0  =  Ba02. 
BaO2=BaO-hO. 


102  -   IXORGANIC   CHEMISTRY 

The  preceding  remodeled  equations  are  ordinary  chemical 
equations;  i.e.  they  are  equations  showing  the  kinds  and 
relative  weights  of  the  interacting  substances. 

The  scope  and  interpretation  of  the  ordinary  chemical 
equation  can  best  be  set  forth  by  a  further  discussion  of  the 
four  kinds  of  chemical  changes. 

(1)  Decomposition.     When  mercuric  oxide  is  heated,  it 
changes   into    mercury   and   oxygen.     This   reaction   is  ex- 
pressed by  the  equation  :  — 

HgO   =    Hg    +     0 

Mercuric       Mercury      Oxygen 
Oxide 

This  equation  may  be  read  in  several  ways:  (a)  Mercuric 
oxide  decomposes  into  mercury  and  oxygen ;  (6)  one  mole- 
cule of  mercuric  oxide  by  decomposition  forms  one  atom  of 
mercury  and  one  atom  of  oxygen;  (c)  216  parts  by  weight 
of  mercuric  oxide  yield  200  parts  by  weight  of  mercury  and 
16  parts  by  weight  of  oxygen  (since  these  are  the  relative 
weights  found  by  experiment  and  reduced  to  the  atomic 
weight  basis);  (d)  mercuric  oxide  equals  mercury  plus 
oxygen. 

(2)  Combination.     When  magnesium  burns  in  air  or  in 
oxygen,  magnesium  oxide  is  formed.     The  equation  for  the 
reaction  is :  — 

Mg      +     O     =     MgO 

Magnesium      Oxygen      Magnesium 
Oxide  ' 

This  equation  may  be  read  as  follows :  (a)  Magnesium  and 
oxygen  combine  to  form  magnesium  oxide ;  (6)  one  atom  of 
magnesium  combines  with  one  atom  of  oxygen  and  forms 
one  molecule  of  magnesium  oxide;  (c)  24  parts  by  weight 
of  magnesium  combine  with  16  parts  by  weight  of  oxygen 
and  yield  40  parts  by  weight  of  magnesium  oxide;  (d)  mag- 
nesium plus  oxygen  equals  magnesium  oxide. 


LAW,   THEORY,   AND   HYPOTHESIS  103 

(3)  Substitution.     When  zinc  and  sulphuric  acid  interact, 
zinc  replaces  the  hydrogen  of  the  sulphuric  acid.     The  equa- 
tion for  this  reaction  is  :  - 

Zn  +  H2S04  =    2  H    +  ZnSO4 

Zinc        Sulphuric      Hydrogen        Zinc 

Acid  Sulphate 

This  equation  may  be  read  as  follows  :  (a)  Zinc  and  sulphuric 
acid  interact  and  form  hydrogen  and  zinc  sulphate  by  the 
substitution  of  zinc  for  the  hydrogen  of  the  acid;  (6)  one 
atom  of  zinc  interacts  with  one  molecule  of  sulphuric  acid 
and  forms  two  atoms  of  hydrogen  and  one  molecule  of  zinc 
sulphate;  (c)  65  parts  by  weight  of  zinc  interact  with  98 
parts  by  weight  of  sulphuric  acid  and  yield  2  parts  by 
weight  of  hydrogen  and  161  parts  by  weight  of  zinc  sulphate; 
(d)  zinc  and  sulphuric  acid  equal  hydrogen  and  zinc 
sulphate. 

(4)  Double  Decomposition.     When  solutions  of  silver  ni- 
trate and  sodium  chloride  are  mixed,  silver  chloride  and 
sodium  nitrate  are  formed.     This  kind  of  chemical  change  is 
called  double  decomposition,  or  metathesis,  because  both  of 
the  original  compounds  undergo  decomposition.     It  is  really 
an  exchange  or  redistribution  of  atoms;    the  original  com- 
pounds undergo  decomposition  while  the  final  compounds 
result  from  a  combination  of  these  parts  on  another  plan. 
Double  decomposition  may  be  regarded  as  the  simultaneous 
occurrence  of  the  other  kinds  of  chemical  change,  for  it  in- 
volves decomposition,  combination,  and  substitution.     How- 
ever, it  is  customary  to  give  this  kind  of  chemical  change 
a  special  name,  owing  to  certain  unique  features  displayed 
by  it.     One  of  these  features  is  the  quite  frequent  forma- 
tion of  an  insoluble  solid  called  a  precipitate.     For  example, 
in  the  reaction  just  cited  the  silver  chloride  is  produced  as 
a  white,  curdy  solid,  almost  insoluble  in  the   final   liquid. 
Double  decomposition  often  results  in  precipitation.     It  is 


104  INORGANIC   CHEMISTRY 

therefore  an  excellent  way  to  make  testa,  and  it  finds  numer- 
ous applications  in  chemical  analysis.  The  equation  for  the 
foregoing  reaction  is  :  — 

AgNO3+  Nad  =  AgCl  +  NaNO3 

Silver          Sodium  Silver  Sodium 

Nitrate        Chloride      Chloride         Nitrate 

This  equation  may  be  read  as  follows:  (a)  Silver  nitrate  and 
sodium  chloride  interact  and  form  silver  chloride  and  sodium 
nitrate  by  double  decomposition;  (6)  one  molecule  of  silver 
nitrate  interacts  with  one  molecule  of  sodium  chloride  and 
forms  one  molecule  of  silver  chloride  and  one  molecule  of 
sodium  nitrate;  (c)  170  parts  by  weight  of  silver  nitrate 
interact  with  58.5  parts  by  weight  of  sodium  chloride  and 
yield  143.5  parts  by  weight  of  silver  chloride  and  85  parts  by 
weight  of  sodium  nitrate;  (d)  silver  nitrate  plus  sodium 
chloride  equals  silver  chloride  plus  sodium  nitrate. 

In  the  foregoing  discussion  of  the  four  kinds  of  chemical 
change  similar  statements  are  designated  by  the  same  letter. 
Let  us  consider  each  lettered  group.  Under  (a)  in  each  case 
nothing  is  said  about  the  physical  conditions  attending  the 
chemical  change  recorded  by  the  equation,  because  such 
accompaniments  are  outside  the  scope  of  ordinary  chemical 
equations.  Thus,  in  the  equations  HgO  =  Hg  +  O  and 
Mg  4-  O  =  MgO  there  is  no  hint  whatever  that  the  mercuric 
oxide  must  be  kept  at  a  high  temperature  or  that  the  mag- 
nesium unites  vigorously  with  oxygen  at  a  relatively  low 
temperature.  Again,  in  the  equations  Zn  +  H2SO4  =  2  H  + 
ZnSO4  and  AgNO8  +  NaCl  =  AgCl  +  NaNO3  there  is  no  sug- 
gestion that  the  sulphuric  acid,  sodium  chloride,  and  silver 
nitrate  must  be  dissolved  in  water,  nor  that  the  zinc  sulphate 
which  is  produced  remains  in  solution  while  the  silver  chloride 
is  precipitated.  Furthermore,  heat  is  often  liberated  in 
chemical  changes,  e.g.  by  the  interaction  of  zinc  and  hydro- 
chloric acid;  but  in  the  ordinary  chemical  equation  as  given 


LAW,   THEORY,   AND   HYPOTHESIS  105 

above  this  fact  is  ignored.  Hence  we  conclude  (1)  that  ordi- 
nary chemical  equations  tell  nothing  about  the  physical 
conditions  (i.e.  temperature,  physical  state,  solution,  etc.) 
under  which  the  chemical  reaction  starts,  proceeds,  and 
ends.  Again,  in  (6)  and  (c)  the  chemical  equations  are  made 
up  of  the  smallest  integral  number  of  atoms  and  molecules 
involved  in  the  chemical  change.  The  equation  is  a  sample, 
so  to  speak,  of  the  vast  number  of  like  changes  which  we 
call  the  chemical  change.  What  determines  the  number  of 
atoms  and  molecules  to  be  incorporated  in  an  equation? 
The  answer  to  this  question  necessitates  the  discussion  of 
several  topics.  First,  equations  are  the  outcome  of  ex- 
periments. They  follow  experiments,  and  are  designed  to 
be  compact,  symbolic  expressions  of  certain  phases  of  a  par- 
ticular chemical  change.  They  tell  at  a  glance  one  part  of 
a  complex  story.  Second,  ordinary  chemical  equations  ex- 
press quantitative  relations.  That  is,  they  not  only  empha- 
size the  fact  that  a  chemical  change  exemplifies  the  law  of 
the  conservation  of  matter,  but  they  also  show  the  propor- 
tions of  the  participating  substances.  For  example,  ex- 
periment shows  that  when  a  given  weight  of  mercuric  oxide 
is  decomposed  into  mercury  and  oxygen,  the  actual  weights 
involved  are  in  the  ratio  of  216  to  200  to  16  respectively. 
Corresponding  values  are  found  for  each  equation.  These 
values,  which  differ  of  course  with  different  equations,  are 
real  and  must  be  known  before  the  particular  equation  can 
be  written  correctly.  In  fact,  they  precede  every  equation, 
although  we  often  overlook  this  fact  in  using  equations. 
Third,  before  a  chemical  equation  can  be  written  certain 
facts  must  be  known.  One  of  these  is  the  composition  of 
each  compound  involved;  i.e.  not  merely  the  per  cent  of  each 
constituent  in  each  compound,  but  the  proportion  of  each 
constituent  in  terms  of  the  atomic  weights  of  the  elements. 
In  other  words,  before .  an  equation  can  be  written  it  is 


106  INORGANIC  CHEMISTRY 

necessary  to  know  the  symbol  and  atomic  weight  of  each 
element  and  the  formula  and  molecular  weight  of  each  com- 
pound. These,  as  previously  stated,  may  be  found  for  the 
present  by  utilizing  the  table  of  atomic  weights  given  in  the 
Appendix,  §  5,  supplemented  by  information  given  in  the 
text.  It  is  also  essential  to  know  the  proportions  in  which 
the  original  substances  (often  called  factors)  interact  and 
in  which  the  final  substances  (sometimes  called  products) 
are  produced. 

We  draw  as  a  second  conclusion  from  this  rather  long  dis- 
cussion of  the  statements  recorded  above  (in  (6)  and  (c))  that 
(2)  an  ordinary  chemical  equation  shows  by  means  of  the 
appropriate  number  of  atoms  and  molecules  not  only  the 
kind  but  the  relative  quantities  of  substances  before  and 
after  a  chemical  change. 

(d)  The  verbal  interpretation  of  ah  ordinary  chemical 
equation  is  very  often  compressed  into  a  form  which  simu- 
lates the  algebraic  equation.  But  ordinary  chemical  equa- 
tions cannot  be  subjected  to  transposition  or  factoring.  In 
a  certain  sense,  therefore,  it  is  anomalous  to  say  that  mercuric 
oxide  equals  mercury  plus  oxygen.  There  will  be  little  or  no 
difficulty,  however,  if  it  is  understood  that  the  ordinary  chem- 
ical equation  is  not  an  algebraic  expression  which  attempts 
to  describe  a  chemical  change  in  all  its  aspects.  Objections 
have  been  raised  to  the  use  of  the  sign  of  equality  (=),  and 
an  arrow  is  sometimes  substituted  for  it,  the  equation  then 
becoming,  for  example,  HgO  ->  Hg  +  O.  Many  equations 
are  used  merely  to  express  the  gravimetric  proportions  in 
which  chemical  reactions  take  place  ;  in  such  equations  we 
shall  use  the  sign  of  equality.  Doubtless  certain  facts  can 
best  be  expressed  by  equations  in  which  the  sign  of  equality 
is  replaced  by  an  arrow.  The  special  style  or  form  of  an 
equation  depends  upon  the  part  of  the  chemical  story  it  is 
designed  to  epitomize.  Later  we  shall  have  occasion  to  use 


LAW,  THEORY,  AND  HYPOTHESIS  107 

other  forms  of  equations.     (See  ionic,  thermal,  and  gas  equa- 
tions, and  equilibrium.) 

We  draw  finally  a  third  and  comprehensive  conclusion ; 
viz.  (3)  ordinary  chemical  equations  are  expressions  show- 
ing by  symbols  and  formulas  the  quantitative  relations  be- 
tween all  the  substances  involved  in  chemical  reactions. 
Each  equation  is  the  outcome  of  experiment,  and  although 
the  equation  contains  signs  used  in  algebra,  a  chemical 
equation  has  none  of  the  properties  of  an  algebraic  equation 
except  equality  between  the  total  weights  on  each  side  of  the 
equation;  the  chemical  equation,  furthermore,  is  limited 
to  a  statement  of  the  chemical  distribution  of  atoms  and 
does  not  reveal  any  facts  about  the  physical  phenomena 
which  invariably  accompany  chemical  changes. 

Making  Equations.  —  It  is  clear  that  the  task  of  making 
a  correct  chemical  equation  is  not  easy.  Several  avenues 
are  open  to  the  beginner.  The  short  equations  can  be  com- 
mitted to  memory  or  worked  out  by  methods  soon  to  be 
outlined;  the  long  ones  can  be  interpreted  by  the  facts 
recorded  in  connection  with  the  experiment  and  then  re- 
ferred to  later,  as  occasion  demands.  The  student  must 
not  forget  that  each  chemical  reaction  has  its  own  equation 
and  that  similarity  of  names  and  of  chemical  changes  does 
not  imply  uniformity  of  equations.  A  word  of  caution 
must  also  be  uttered  against  attempts  to  write  an  equation 
by  mere  guess  work  and  then  expect  the  facts  to  coincide 
with  this  pseudo-equation. 

One  method  of  working  out  simple  chemical  equations 
will  be  clear  from  the  following  cases :  (a)  When  magnesium 
is  heated  in  oxygen  (or  in  air),  the  ratio  by  weight  in  which 
the  two  elements  combine  is  3:2.  This  result  is  expressed 
in  terms  of  the  atomic  weights  of  the  elements  involved. 
Let  y  equal  the  number  of  atomic  weights  of  magnesium 


108  INORGANIC  CHEMISTRY 

and  z  the  number  of  atomic  weights  of  oxygen.  Then  we 
can  write  a  preliminary  equation  thus :  — 

y  x  at.  wt.  of  magnesium  :  z  x  at.  wt.  of  oxygen  =  3:2. 

The  atomic  weight  of  magnesium  is  found  by  the  table  on 
the  back  cover  to  be  24  and  that  of  oxygen  to  be  16. 
Substituting  these  values,  we  have  the  equation  :  — 

2/x24:zxl6  =  3:2 

By  inspection  y  =  z,  and  the  simplest  value  of  each  is  1. 
Now  the  symbol  Mg  expresses  24  parts  of  magnesium  and 
the  symbol  O  expresses  16  parts  of  oxygen.  Therefore 
Mg  and  O  are  the  symbols  representing  the  smallest  number 
of  atoms  equivalent  arithmetically  to  the  ratio  (3  :  2)  found 
by  experiment.  The  formula  of  the  product  formed  by 
their  combination  is  therefore  MgO,  and  the  simplest  equa- 
tion expressing  the  chemical  change  is  — 

Mg       +     0     =     MgO 

Magnesium        Oxygen      Magnesium 
Oxide 

(6)  Again,  suppose  we  wish  to  find  the  simplest  equation 
for  the  reaction  between  hydrogen  and  oxygen  in  the  for- 
mation of  water.  Experiment  shows  that  hydrogen  and 
oxygen  combine  in  the  ratio  1  : 8  by  weight.  Pursuing  the 
same  line  of  argument  as  above,  let  y  =  the  number  of  atomic 
weights  of  hydrogen,  and  z  that  of  oxygen.  The  preliminary 
equation  is  :  — 

y  X  at.  wt.  of  hydrogen  :  z  x  at.  wt.  of  oxygen  =  1:8 

The  atomic  weight  of  hydrogen  is  1  and  of  oxygen  is  16. 
The  equation  now  becomes  — 

y  X  1 :  z  X  16  •»  1 :  8 

By  inspection,  y  =  2  z,  and  the  simplest  values  are  y  =  2 
and  z  =  1.  Now  the  symbol  H  stands  for  1  part  of  hydrogen, 


LAW,   THEORY,   AND   HYPOTHESIS  109 

and  O  for  16  parts  of  oxygen.  Therefore,  2  H  and  O  are 
the  symbols  representing  the  smallest  number  of  atoms 
equivalent  arithmetically  to  the  ratio  (1:8)  found  by  ex- 
periment. The  formula  of  the  product  of  their  combina- 
tion is  H2O,  and  the  simplest  equation  is  — 

2  H  +  0  =  H2O 

All  equations  are  not  equally  simple,  but  by  a  similar 
argument  based  upon  the  facts  found  by  experiment  many 
simple  equations  may  be  developed. 

Another  method  is  often  possible.  When  we  know  the 
factors  and  products  of  a  reaction,  we  can  find  their  symbols 
or  formulas  in  the  book,  construct  a  preliminary  equation, 
and  then  balance  the  equation;  i.e.  select  the  proper  coeffi- 
cients, subscripts,  or  both,  so  that  there  shall  be  an  equal 
number  of  atoms  of  each  element  on  both  sides  of  the  equa- 
tion. An  example  will  make  this  method  clear.  When 
phosphorus  burns  in  oxygen,  phosphorus  pentoxide  is  formed. 
The  preliminary  equation  is  — 

P  +  O  =  P205 

Here  it  is  evident  that  to  balance  the  equation  we  need  2  P 
and  5  O  on  the  left.  Hence  the  final  equation  is  — 


Again,  when  zinc  and  hydrochloric  acid  interact,  hydrogen 
and  zinc  chloride  are  formed.  The  preliminary  equation 
made  from  the  symbols  and  formulas  is  — 

Zn  -|-  HC1  =  H  4-  ZnCl2 

By  inspection,  it  is  evident  that  two  atoms  of  chlorine  are 
on  the  right  and  only  one  on  the  left.  To  obtain  C12  it  is 
necessary  to  write  2  HC1.  But  2  HC1  means  not  only  2  Cl 
but  2  H.  Hence  the  equation  becomes  — 

Zn-f  2HCl  =  2H  +  ZnCla 


110  INORGANIC   CHEMISTRY 

A  final  inspection  shows  that  an  equal  number  of  atoms  of 
each  element  is  on  both  sides  of  the  equation. 

Many  equations  may  be  written  by  applying  these  methods 
to  the  facts  found  by  experiment  (see  exercises  at  the  end  of 
this  chapter). 

Equations  for  Preceding  Reactions.  — The  equations  cor- 
responding to  many  reactions  already  discussed  may  ap- 
propriately, be  collected  here,  partly  for  their  retrospective 
value  and  partly  for  future  use.  The  equation  for  the 
preparation  of  oxygen  from  mercuric  oxide  is  — 

HgO  =     Hg    +     O 

Mercuric      Mercury      Oxygen 
Oxide 

When  sulphur,  carbon,  magnesium,  and  iron  are  burned  in 
air  (or  in  oxygen),  the  equations  are  — 

S      +    2O    =    S02 

Sulphur       Oxygen       Sulphur 
Dioxide 

c    +  20  =  co2 

Carbon        Oxygen        Carbon 
Dioxide 

Mg      +     O  .  =     MgO 

Magnesium        Oxygen       Magnesium 
Oxide 

3Fe+    4O    =     Fe3O4 

Iron        Oxygen          Magnetic 
Iron  Oxide 

The  equation  for  the  preparation  of  hydrogen  from  zinc  and 
hydrochloric  acid  is  — 

Zn  +      2HC1     =     2H     +  ZnCl2 

Zinc       Hydrochloric       Hydrogen  Zinc 

Acid  Chloride 

When  hydrogen  burns  or  when  a  mixture  of  hydrogen  and 
oxygen  is  exploded,  the  equation  is  — 


LAW,   THEORY,  AND  HYPOTHESIS  111 

2  H      +      O     =  H2O 

Hydrogen       Oxygen       Water 

The  equation  for  the  reaction  in  determining  the  gravimetric 
composition  of  water  is  — 

CuO  +     2H     =  H2O  +    Cu 

Copper       Hydrogen       Water       Copper 
Oxide 

The  equation  for  the  decomposition  of  potassium  chlorate  is  — 
KC103  =      KC1     +30 

Potassium         Potassium       Oxygen 
Chlorate  Chloride 

The  interaction  of  sodium  and  water  is  represented  thus:  — 
Na    +  H2O  =       H       +    NaOH 

Sodium       Water        Hydrogen  Sodium 

Hydroxide 

When  phosphorus  burns  in  air  (or  oxygen),  the  equation  is  — 
2P      +    5O    =      P2O5 

Phosphorus       Oxygen       Phosphorus 
Pentoxide 

Calculations  based  on  Equations.  —  Since  equations  are 
expressions  of  the  relative  quantities  of  the  substances  in- 
volved in  chemical  reactions,  it  is  possible  to  solve  many 
arithmetical  problems  arising  from  reactions.  An  equation 
states  the  proportions  which  participate  chemically  in  a 
reaction.  Obviously,  any  convenient  weights  of  zinc  and 
sulphuric  acid  might  be  brought  together,  but  the  pro- 
portions according  to  which  the  factors  react  and  the  prod- 
ucts are  formed  are  always  expressed  by  the  equation  — 

Zn  +  H2SO4  =     2  H     +  ZnSO4 

Zinc        Sulphuric        Hydrogen  Zinc 

Acid  Sulphate 

65  98  2  161 

If  zinc  and  sulphuric  acid  are  brought  together  in  any  other 
proportion,  a  part  of  one  or  the  other  will  be  left  over  unused. 


112  INORGANIC  CHEMISTRY 

The  equation  above  means  that  zinc  and  sulphuric  acid 
always  interact  in  the  proportion  of  65  and  98,  and  produce 
hydrogen  and  zinc  sulphate  in  the  proportion  of  2  and  161. 
We  may  read  grams,  ounces,  kilograms,  or  any  other  unit 
in  connection  with  these  numbers,  but  the  same  unit  must 
be  used  throughout  the  calculations.  Therefore,  if  we  know 
the  actual  weight  of  one  substance  participating  in  a  reaction, 
all  other  weights  involved  can  be  readily  calculated.  Sup- 
pose 45  gm.  of  zinc  interact  with  sulphuric  acid;  the  weights 
of  (a)  acid  required,  (6)  hydrogen  formed,  and  (c)  zinc 
sulphate  produced  are  calculated  as  follows  :  — 

(1)  Write  the  chemical  equation  for  the  reaction,  thus  :  — 

Zn  +  H2S04  =  2  H  +  ZnSO4 

(2)  Place  under  each  term  of  the  equation  its  atomic  or 
molecular  weight,1  as  the  case  may  be,  thus  :  — 


65  98  2  161 

(3)  Place  above  the  proper  terms  the  known  weight  and 
required  weight  (i.e.  X,  Y,  Z,  etc.)  involved  in  the  problem, 
thus:  — 

45  X  Y  Z 


65  98  2  161 

(4)  State  in  the  form  of  a  proportion  the  four  terms  in- 
volved, remembering  that  the  known  and  required  weights 
are  in  the  same  ratio  as  the  atomic  and  molecular  weights. 
Thus,  the  three  proportions  in  the  given  problem  are:  — 

(a)  45  :  X  :  :  65  :  98;    X  =   67.8  gm.  sulphuric  acid. 
(6)  45  :  Y  :  :  65  :  2  ;       Y  =    1.38  gm.  hydrogen. 
(c)  45  :  Z  :  :  65  :  161  ;    Z  =  111.4  gm.  zinc  sulphate. 
Similar  problems  can  be  solved  by  this  method. 

1  The  atomic  weights  are  given  in  the  table  on  the  inside  of  the  back 
cover.  Molecular  weights  are  obtained  by  adding  the  proper  atomic  weights. 


LAW,  THEORY,  AND  HYPOTHESIS  113 

PROBLEMS  AND  EXERCISES 

1.  Calculate  the  percentage  composition  of  (a)  water  (H20), 
(6)   zinc   oxide    (ZnO),    (c)   lead   carbonate    (PbCO3),    (d)  sodium 
chlorate  (NaC103),  (e]  barium  oxide  (BaO),  (/)  calcium  carbonate 
(CaC03). 

2.  Calculate   the  percentage  composition  of    (a)   copper  sul- 
phate (CuSO4),  (6)  barium  chloride  (BaCl2),  (c)  manganese  dioxide, 
(d)  calcium  oxide,  (e)  sodium  hydroxide,  (/)  potassium  hydroxide, 
(g)   sodium  carbonate   (Na2C03),    (h)  potassium  nitrate   (KNO3), 
(t)  mercuric  oxide  (HgO). 

3.  Show  that  the  following  sets  of  compounds  illustrate  the 
law  of  multiple  proportions :   (a)  H  =  11.11  per  cent  and  O  =  88.88 
per   cent,    H  =  5.882   and   0  =  94.117;     (6)  Sn  =  62.63  and  Cl  = 
37.37,  Sn  =  45.49  and  Cl  =  54.41. 

4.  Calculate  the  formula  of  the  compounds  which  have  the 
indicated  composition :  (a)  Na  =  60.68,  Cl  =  39.31 ;   (6)  Ca  =  29.41 ; 
8  =  23.52,  0  =  47.05;  (c)  C  =  27.27,  0  =  72.72;  (d)  As  =  75.8,  O  = 
24.2;  (e)  N  =  82.35,  H=  17.63. 

5.  Calculate  the  formula  of  the  compounds  which  have  the  fol- 
lowing composition :    (a)  Si  =  19.5,  C  =  66.62,  H  =  13.88 ;    (6)  Pb  = 
86.6,  S  =  13.4 ;    (c)  N  =  26.17,  H  =  7.48,  Cl  =  66.35. 

6.  How  much  oxygen  can  be  prepared  from  (a)  70  gm.  of  mer- 
curic oxide  ;    (6)   17  gm.  of  potassium  chlorate  ? 

7.  How  much  potassium  chloride  will  remain  after  82.5  1.  of 
oxygen  (at  0°  C.  and  760  mm.)  have  been  obtained  from  potassium 
chlorate  ? 

8.  Interpret  the  following :     H,  H2,  2  H,  H2O,'  2  H20,  H2O2, 
NaOH,  Ca(OH)2,  A1(NO»)8,  ZnCl2,  2  ZnSO4,  3  Fe3O4,  5  P2O6. 

9.  Interpret  the  following  :  (a)  PbO2  =  PbO  +  O  ;  (6)  Cu  +  O  = 
CuO;    (c)Zn  +  2NaOH=Na2ZnO2  +  2H;  (d)  Ba(NO3)2+  K2SO4  = 
BaSO4  +  2  KN03. 

10.  From  the  equation  KC103  =  KC1  +  3  O,  calculate   (a)  the 
weight  of  potassium  chloride  when  30  gm.  of  oxygen  are  liberated ; 
(6)  the  weight  of  chlorine  the  potassium  chloride  will  yield,  if  10 
gm.  of  potassium  chlorate  are  decomposed ;      (c)  the  volume  of 
oxygen  liberated  (at  0°  C.  and  760  mm.)  from  35  gm.  of  potassium 
chlorate  containing  4.5  gm.  of  potassium  chloride. 

11.  Iron  and  sulpur  unite  in  the  ratio    of   7:4.     Write    the 
equation  for  the  reaction.     What  weight  of  the  product  will  be 
formed  from  (a)  20  gm.  of  iron  and  (6)  20  gm,  of  sulphur  ? 


114  INORGANIC  CHEMISTRY 

12.  What  weight  of  (a)  zinc  and  (6)  hydrochloric  acid  are  needed 
to  produce  17.5  gm.  of  hydrogen?     17.5  1.  (at  0°  C.  and  760  mm.)  ? 

13.  How  much  ferric  oxide  will  yield  2.5  gm.  of  iron? 

14.  How  much   silver  chloride  can  be   produced  from   silver 
nitrate  and  22.8  gm.  of  crystallized  barium  chloride  (BaCl2  .  2  H2O)  ? 

15.  Calculate  the  total  weight  of  water  that  can  be  obtained 
from  a  metric  ton  of  crystallized  calcium  sulphate. 

16.  A  balloon  holds  150  kg.  of  hydrogen.     How  much  (a)  zinc 
and  (6)  sulphuric  acid  are  needed  to  produce  the  gas  ? 

17.  If  water  and  10  gm.  of  sodium  interact,  calculate  the  weight 
of  each  product. 

18.  In  the  reduction  of  copper  oxide  by  hydrogen  2.52  gm.  of 
the  solid  product  resulted.      What  weights  of  copper  oxide  and 
hydrogen  were  used  ? 

19.  A  liter  of  oxygen  (at  0°  C.  and  760  mm.)  was  transformed 
by    phosphorus   into    phosphorus    pentoxide.     How   many  grams 
of  phosphorus  and  of  phosphorus  pentoxide  were  involved  ? 

20.  A  lump  of  carbon  weighing  10  gm.  is  burned  in  air.     What 
(a)  volume  and  (6)  weight  of  carbon  dioxide  is  formed? 

21.  What  weight  of  iron  oxide   (Fe304)  is  formed  by  burning 
a  metric  ton  of  iron  in  oxygen  ?     What  volume  of  oxygen  (at  0°  C. 
and  760  mm.)  is  used? 

22.  A  lump  of  sulphur  weighing  12  gm.  is  burned  in  air.     Calcu- 
late (a)  the  weight  and  (6)  the  volume  of  oxygen  needed  and  sul- 
phur dioxide  formed.     If  air  contains  21  per  cent  of  oxygen  (by 
volume),  what  volume  of  air  is  used? 

23.  Calculate  the  weight  of  oxygen  needed  to  burn  33  gm.  of 
magnesium  containing  11  per  cent  of  impurities. 

24.  A  liter  of  hydrogen  (at  0°  C.  and  760  mm.)  is  produced 
by  the  interaction  of  aluminium  and  sodium  hydroxide.     What 
weight  of  the  other  substances  are  involved  ? 

25.  Write  the  equation  for  the  combination  of  sulphur  and 
oxygen  when  they  unite  in  the  ratio  (by  weight)  (a)  1 :  1  and  (6)2:3; 
for  carbon  and  oxygen  in  the  ratio  (a)  3  :  8  and  (6)  3  :  4 ;   for  nitro- 
gen and  hydrogen  in  the  ratio  4.66 :  1 ;    for  nitrogen  and  oxygen 
in  the  ratios  7:4;   3.5 :  8 ;   3.5 :  6 ;   3.5 :  10. 

26.  Write  equations  for  the  following  reactions  :   (a)  Magnesium 
and    sulphuric    acid    form    magnesium    sulphate    and    hydrogen. 
(6)  Zinc  sulphate  and  barium  nitrate  form  barium  sulphate  and 
zinc  nitrate,     (c)  Strontium  carbonate  and  hydrochloric  acid  form 
strontium  chloride,  water,  and  carbon  dioxide. 


CHAPTER  VIII 
The  Atmosphere  —  Argon  and  Related  Elements  —  Nitrogen 

THE  atmosphere  is  the  great  envelope  of  gas  surrounding 
the  earth.  It  extends  into  space  to  an  estimated  height 
of  fifty  to  two  hundred  miles.  We  live  at  the  bottom  of 
this  vast  ocean  of  air,  as  it  is  often  called. 

Aristotle  (384-322  B.C.)  regarded  air  as  one  of  the  four  elementary 
principles  whose  combinations  made  up  all  substances  in  the  universe. 
The  other  three  were  earth,  fire,  and  water.  He  taught  that  air  pos- 
sesses two  fundamental  properties  —  heat  and  dampness.  The  early 
chemists  used  the  word  air  in  the  sense  in  which  the  word  gas  is  now 
employed.  Thus,  we  have  already  learned  that  hydrogen  was  first 
called  inflammable  air. 

The  terms  atmosphere  and  air  are  often  used  interchange- 
ably, though  by  air  we  usually  mean  a  limited  portion  of 
the  atmosphere.  Many  skillful  chemists  have  studied  the 
action  of  air  on  living  things,  its  relation  to  combustion,  the 
effect  of  its  weight,  its  composition,  and  its  varied  properties. 
Their  work  has  contributed  many  fundamental  facts  to 
.science. 

General  Properties  of  the  Atmosphere.  — Air  has  weight. 
We  often  use  the  expression  "  light  as  air."  But  a  cubic 
foot  of  air  weighs  1.28  oz.  and  a  room  40  X  50  X  25  ft.  con- 
tains about  two  tons  of  air.  A  liter  of  dry,  normal  air  at 
0°  C.  and  760  mm.  weighs  1.293  gm.  The  total  weight  of 
the  atmosphere  has  been  estimated  to  be  five  thousand 
millions  of  millions  of  tons.  The  enormous  mass  resting 
upon  the  earth  exerts  a  pressure  which  is  about  fifteen  pounds 
on  every  square  inch.  The  amount  of  pressure  upon  a 

115 


116  INORGANIC  CHEMISTRY 

square  inch  is  called  "  an  atmosphere,"  and  it  is  sometimes 
used  as  a  unit  of  pressure;  e.g.  three  atmospheres  means 
a  pressure  of  forty-five  pounds  per  square  inch.  It  is  atmos- 
pheric pressure  which  causes  water  to  rise  in  pumps  and 
flow  through  siphons.  Atmospheric  pressure  is  exerted  in  all 
directions  and  is  variable.  It  is  measured  by  the  barometer. 
The  normal  or  standard  pressure  of  the  atmosphere,  as  al- 
ready stated,  is  equal  to  the  pressure  of  a  column  of  mercury 
which  is  760  mm.  (or  29.92  in.)  high.  In  very  accurate  ex- 
periments certain  mathematical  corrections  must  be  made 
in  the  height  of  the  column  as  read  on  the  barometer  scale, 
but  in  ordinary  work  it  is  necessary  to  know  the  height  only 
of  the  mercury  column  in  order  to  know  the  pressure.  The 
pressure  of  the  atmosphere  varies  as  the  height  and  com- 
position of  the  atmosphere  vary,  and  the  barometer  changes 
accordingly. 

Ingredients  of  the  Atmosphere.  —  The  atmosphere  is  a 
mixture  of  several  gases.  -But  since  this  mixture  always 
contains  approximately  78  parts  of  nitrogen  and  21  parts 
of  oxygen  by  volume,  we  often  speak  of  air  as  consisting 
solely  of  these  two  gases.  Besides  this  large  proportion 
of  oxygen  and  nitrogen,  the  air  always  contains  small  and 
variable  proportions  of  water  vapor  and  carbon  dioxide  gas. 
In  addition  to  these  four  ingredients,  air  always  contains 
the  gas  argon  (and  the  related  inert  gases),  and  usually 
very  small  proportions  of  ozone,  hydrogen,  hydrogen  di- 
oxide, compounds  related  to  ammonia  and  nitric  acid,  dust, 
and  germs.  The  composition  varies  but  slightly  in  dif- 
ferent localities.  Near  the  city  air  may  contain  a  relatively 
larger  proportion  of  dust,  ammonia,  sulphur  compounds, 
and  acids ;  in  the  country  the  proportion  of  ozone  is  rela- 
tively large ;  over  the  ocean  and  near  the  seacoast  the  air 
contains  salt. 


THE  ATMOSPHERE  117 

General  Properties  of  Nitrogen.  — The  chemical  element, 
nitrogen,  which  constitutes  about  78  per  cent  of  the  at- 
mosphere (by  volume),  is  a  colorless  gas,  and  has  no  taste 
or  odor.  It  is  somewhat  lighter  than  air,  and  is  slightly 
soluble  in  water.  In  many  respects  it  differs  markedly 
from  oxygen.  Thus,  it  will  not  support  combustion,  neither 
will  it  burn,  nor  sustain  life.  Animals  die  if  left  in  nitrogen. 
Nitrogen  is  not  poisonous,  for  the  air  we  breathe  contains 
a  large  proportion  of  nitrogen.  Its  function  in  the  at- 
mosphere is  to  dilute  the  oxygen.  It  is  an  inert  element 
compared  with  many  others,  although  it  combines  directly 
with  oxygen  and  a  few  other  elements.  At  the  ordinary 
temperature  it  is  chemically  indifferent,  but  at  high  tempera- 
tures it  is  quite  active  (see  pages  129,  212,  217). 

The  fact  that  a  candle  flame  quickly  goes  out  and  a  mouse  soon 
dies  in  nitrogen  was  first  observed  by  Rutherford,  a  Scottish  physician, 
who  discovered  the  gas  in  1772.  Soon  after,  Lavoisier  showed  the 
true  relation  of  nitrogen  to  the  atmosphere.  To  emphasize  the  ina- 
bility of  the  gas  to  support  life,  he  called  the  new  gas  azote,  the  name 
now  used  for  it  by  French  chemists. 

Oxygen  and  Nitrogen  in  the  Atmosphere.  — The  chemical 
activity  of  the  atmosphere  is  due  to  the  free  oxygen  it  con- 
tains, as  we  have  already  learned  in  studying  oxygen.  If 
the  air  were  largely  oxygen,  rusting  and  decay  would  proceed 
with  astonishing  rapidity,  and  fires  once  started  would  burn 
with  great  violence.  On  the  other  hand,  nitrogen  is  chemi- 
cally inactive,  and  if  the  air  contained  much  more  than  the 
normal  amount,  the  chemical  action  of  oxygen  would  be 
slower.  Oxygen  alone  is  too  active,  while  nitrogen  alone 
is  rather  inactive.  To  be  serviceable  to  man,  oxygen  must 
be  diluted  with  nitrogen,  while  nitrogen  must  be  accom- 
panied by  a  small  proportion  of  oxygen. 

The  presence  of  oxygen  and  nitrogen  in  the  atmosphere  and  the 
functions  of  the  two  gases  were  first  clearly  explained  by  Lavoisier 


118  INORGANIC  CHEMISTRY 

in  1777,  though   many   others  —  Boyle,  Priestley,  Rutherford,   and 
Scheele  —  helped  solve  the  problem. 

Composition  of  the  Atmosphere.  —  Samples  of  air  from 
various  parts  of  the  globe  show  a  remarkable  uniformity 
of  composition.  For  many  years  it  was  believed  that  pure 
air  consisted  solely  of  oxygen  and  nitrogen.  But  in  1895  it 
was  found  that  nearly  2  per  cent  (by  weight)  of  the  gas 
hitherto  called  nitrogen  is  argon.  (See  Argon,  below.)  Ac- 
cording to  the  most  recent  results,  the  following  is  — 

THE  COMPOSITION  OF  PURE  NORMAL,  DRY  AIR 


INGREDIENT 

PERCENTAGE 

By  Volume 

By  Weight 

Nitrogen      ...          .     •     .     . 

78.122 
20.941 
0.937 

75.539 
23.024 
1.437 

The  composition  of  the  atmosphere  was  studied  by  Priestley, 
but  his  results  were  conflicting.  Cavendish,  in  1781,  was  the  first  to 
ehow  that  the  proportion  of  oxygen  and  nitrogen  in  air  is  nearly 
constant.  Since  his  time  this  result  has  been  confirmed  by  many 
chemists,  especially  by  Bunsen.  In  recent  years  the  composition  of 
the  atmosphere  and  the  properties  of  its  inert  components  have  been 
assiduously  studied  by  Ramsay  and  others. 

The  Volumetric  Composition  of  Air  can  be  found  in  the 
laboratory  by  introducing  a  known  volume  of  pure  air  into 
a  eudiometer  and  exploding  it  with  a  known  and  sufficient 
volume  of  hydrogen.  The  nitrogen  does  not  participate 
in  the  chemical  change,  but  all  the  oxygen  in  the  air  combines 
'with  twice  its  volume  of  hydrogen,  forming  a  minute  quantity 
of  water;  hence  one  third  of  the  diminution  in  volume  is 
the  volume  of  oxygen  in  the  air.  The  difference  between 


THE  ATMOSPHERE 


119 


the  volume  of  oxygen  found  and  the  original  volume  of  air 
is  the  volume  of  the  other  constituents  —  chiefly  nitrogen. 
An  illustration  will  make  this  experiment  clear.  Suppose 
(1)  we  mix  and  explode  100  cc.  of  air  and  50  cc.  of  hydrogen, 
or  150  cc.  in  all,  and  (2)  the  residue  measures  87  cc.  Now, 
150  —  87  =  63,  hence  63  cc.  of  the  total  volume  of  the  mix- 
ture combined  to  form  water  (the  63  consisting  of  all  the  oxy- 
gen and  part  of  the  hydrogen).  But  one  third  of  the  63  cc.  is 
the  oxygen  which  was  in  the  original  volume  of  air,  because 
oxygen  and  hydrogen  unite  volumetrically  in  the  ratio  of  1 
to  2.  Hence,  63  -f-  3  =  21,  the  volume  of  oxygen  in  100  cc. 
of  air.  The  remainder,  79  cc.,  is  nitro- 
gen, argon,  and  other  gases. 

Another  Method,  often  used  to  determine 
the  volumetric  composition  of  the  air,  is  based 
on  the  fact  that  phosphorus  combines  slowly 
with  oxygen,  even  at  the  ordinary  tempera- 
ture. The  operation  is  performed  by  insert- 
ing a  piece  of  phosphorus  (see  Figure  13  a) 
into  a  graduated  glass  tube  containing  a  meas- 
ured volume  of  air.  White  fumes  indicate 
immediate  action.  These  fumes  are  solid  par- 
ticles of  phosphorus  pentoxide.  They  soon 
dissolve  in  the  water,  which  rises  higher  in 
the  tube,  as  the  oxygen  combines  with  the 
phosphorus.  In  a  few  hours  the  phosphorus 
is  removed,  and  the  volume  of  gas  is  read. 
The  difference  between  the  first  and  last  vol-  ( 
umes  is  oxygen.  The  gas  remaining  in  the 
tube  is  a  mixture  of  nitrogen  and  argon.  In 
performing  this  experiment  unusual  care 
must  be  taken  not  to  touch  the  phosphorus 
with  dry  hands.  Both  gas  volumes  should 
be  corrected  for  pressure,  temperature,  and  aqueous  vapor. 

The  Gravimetric  Composition  of  Air,  as  already  stated, 
is :  — 


FIG.  13  a.  —  Apparatus 
for  determining  the 
volumetric  composi- 
tion of  air. 


120  INORGANIC  CHEMISTRY 

Oxygen 23.024  parts  by  weight, 

Nitrogen 75.539  parts  by  weight, 

Argon 1.437  parts  by  weight. 

Dumas  and  Boussingault  in  1841  made  the  first  accurate  deter- 
mination of  the  gravimetric  composition  of  the  air.  They  passed 
pure  air  through  a  weighed  tube  containing  copper,  and  arranged 
so  that  heat  could  be  applied.  The  oxygen  of  the  air  combined 
with  the  copper,  while  the  nitrogen  passed  on  into  a  weighed  globe. 
Both  tube  and  globe  increased  in  weight.  The  increase  in  the  tube 
was  the  weight  of  the  oxygen,  while  the  increase  in  the  globe  was 
the  weight  of  the  nitrogen. 

Water  Vapor  in  the  Atmosphere.  —  Water  vapor  is  always 
present  in  the  atmosphere,  owing  to  constant  evaporation 
from  the  ocean,  other  bodies  of  water,  and  the  soil.  The 
total  amount  is  large,  though  variable.  A  given  volume  of 
air  at  a  given  temperature  will  absorb  a  definite  volume  of 
water  vapor  and  no  more.  The  amount  absorbed  depends 
largely  upon  the  temperature.  Air  containing  its  maximum 
amount  of  water  vapor  at  a  given  temperature  is  said  to  be 
saturated  at  that  temperature,  or  to  contain  100  per  cent 
of  water  vapor.  The  saturation  point  is  also  called  the 
dew-point.  On  a  pleasant  day,  the  relative  humidity,  i.e. 
the  relative  amount  of  water  vapor  present,  may  vary  from 
30  to  90  per  cent,  the  average  being  about  50  per  cent, 
though  it  varies  with  the  locality.  Warm  air  holds  more 
vapor  than  cool  air.  The  amount  of  water  vapor  in  the  air 
has  a  marked  influence  on  the  physical  condition  of  man. 
The  depressing  weather  during  "  dog  days  "  is  due  to  the 
high  relative  humidity  of  the  air,  which  sometimes  reaches 
nearly  100  per  cent.  The  specialized  forms  of  life,  both 
animal  and  vegetable,  found  in  deserts  are  largely  due  to 
the  dry  air.  The  languor  felt  in  a  "  close  "  room  or  crowded 
hall  is  partly  caused  by  the  excess  of  water  vapor  in  the 
"  bad  "  air.  The  presence  of  water  vapor  in  the  air  is  shown 


THE  ATMOSPHERE  121 

by  the  moisture  which  collects  on  the  outside  of  a  vessel 
containing  cold  water,  such  as  a  pitcher  of  iced  water.  The 
moisture  comes  from  the  air  around  the  vessel.  For  a 
similar  reason,  water  pipes  in  a  cellar  and  the  cellar  walls 
themselves  are  moist  in  summer.  The  deliquescence  of 
calcium  chloride,  common  salt,  and  other  substances  like- 
wise reveals  the  presence  of  water  vapor  in  the  air.  (See 
Deliquescence.)  When  the  temperature  of  the  air  falls 
sufficiently,  the  water  vapor  condenses  and  is  deposited  in 
the  form  of  dew,  rain,  fog,  mist,  frost,  snow,  sleet,  or  hail. 
The  clouds  are  masses  of  water  vapor  which  has  been  con- 
densed by  the  cold  of  the  upper  air. ' 

Carbon  Dioxide  in  the  Atmosphere.  — Carbon  dioxide 
is  one  product  of  the  respiration  of  animals  and  of  the 
combustion  and  decay  of  organic  substances.  By  these 
processes  an  immense  quantity  of  carbon  dioxide  is  being 
constantly  poured  into  the  atmosphere.  The  proportion  in 
the  atmosphere  is  variable,  though  not  between  such  wide 
limits  as  the  water  vapor.  The  proportion  in  normal  air 
is  3  to  4  parts  in  10,000  parts  of  air.  Over  the  ocean  the 
proportion  is  smaller,  but  in  the  air  of  cities  it  is  greater. 
In  crowded  rooms  the  proportion  is  often  as  high  as  33 
parts  in  10,000.  The  proportion  of  carbon  dioxide  in  the 
atmosphere  as  a  whole  is  practically  constant,  largely  owing 
to  the  fact  that  this  gas  is  an  essential  food  of  plants.  (See 
Carbon  Dioxide.)  The  presence  of  carbon  dioxide  in  the 
air  is  detected  by  calcium  hydroxide.  If  a  solution  of 
calcium  hydroxide  is  exposed  to  the  air,  the  carbon  dioxide 
interacts  with  the  calcium  hydroxide,  forming  a  thin,  white 
crust  of  insoluble  calcium  carbonate  on  the  surface  of  the 
liquid.  If  air  is  drawn  through  calcium  hydroxide,  the  liquid 
becomes  milky,  because  the  fine  particles  of  calcium  car- 
bonate remain  temporarily  suspended  in  the  liquid.  The 


122  INORGANIC  CHEMISTRY 

purity  of  air  is  often  determined  by  finding  out  what  pro- 
portion of  carbon  dioxide  it  contains.  If  a  known  volume 
of  dry  air  is  drawn  through  a  known  weight  of  calcium 
hydroxide  or  similar  liquid,  the  increase  in  weight  will  be 
the  weight  of  carbon  dioxide  in  the  volume  of  air  used. 
The  equation  for  the  interaction  of  carbon  dioxide  and 
calcium  hydroxide  is  :  — 

CO8    4-     Ca(OH)2    =    CaC03     +   H2O 

Carbon  Calcium  Calcium  Water 

Dioxide  Hydroxide  Carbonate 

The  different  gases  in  the  atmosphere  are.  not  arranged 
in  layers  according  to  their  densities.  They  are  in  constant 
circulation.  (See  Diffusion.)  Hence  carbon  dioxide,  though 
heavier  than  oxygen  and  nitrogen  (volume  for  volume), 
does  not  remain  nearest  the  ground,  but  is  distributed 
through  the  air.  In  a  few  exceptional  localities,  carbon 
dioxide  arises  from  volcanic  openings  faster  than  it  can 
diffuse,  and  fills  the  cave  or  adjacent  valley. 

Argon  (A)  and  Related  Elements.  —  Argon,  as  already 
stated,  is  an  essential  and  constant  component  of  the  atmos- 
phere. Argon  was  discovered  in  1895  by  Rayleigh  and  Ram- 
say. Rayleigh  found  that  nitrogen  extracted  from  air  had 
a  greater  weight  than  an  equal  volume  of  nitrogen  obtained 
from  compounds  of  nitrogen.  Consequently,  they  believed 
that  the  nitrogen  from  air  contained  another  gas  hitherto 
overlooked.  A  series  of  elaborate  experiments  showed 
that  after  the  oxygen  and  nitrogen  were  removed  from 
purified  air,  there  still  remained  a  small  quantity  of  a  new 
gas,  which  they  called  argon.  It  may  be  obtained  (1)  by 
passing  pure  air  over  heated  copper  to  remove  the  oxygen, 
and  then  the  remaining  gas  over  heated  .magnesium  or 
calcium  to  remove  the  nitrogen;  or  (2)  by  passing  electric 
sparks  through  a  mixture  of  air  and  oxygen,  and  removing 


SIR    WILLIAM    RAMSAY 


THE  ATMOSPHERE  123 

the  compound  of  oxygen  and  nitrogen  as  fast  as  it  is  formed. 
The  latter  method  is  a  repetition  of  the  one  used  by  Cavendish 
when  he  determined  the  composition  of  air,  and  he  would 
have  no  doubt  discovered  argon  had  he  continued  his 
investigations.  As  stated  above,  the  proportion  in  the 
atmosphere  is  .937  per  cent  by  volume  and  1.437  per  cent 
by  weight. 

Argon  is  a  colorless,  odorless  gas.  It  dissolves  in  water  to 
the  extent  of  about  4  volumes  in  100.  One  liter  of  argon  at 
0°  C.  and  760  mm.  weighs  1.7809  gm.  (compare  with  nitrogen, 
page  129).  Liquid  argon  boils  at  —186°  C.  and  solid  argon 
melts  at  —189.5°  C.  A  conspicuous  property  of  argon  is 
its  lack  of  chemical  activity.  No  compounds  of  this  element 
have  as  yet  been  prepared  or  discovered.  The  name  argon 
is  happily  chosen,  being  derived  from  Greek  words  signify- 
ing inert. 

Helium  (He),  neon  (Ne),  krypton  (Kr),  and  xenon  (Xe) 
are  inert  gases  discovered  by  Ramsay  subsequently  to  argon. 
They  constitute  an  exceedingly  minute  proportion  of  the 
atmosphere.  Like  argon,  they  do  not  form  compounds. 
Their  proportions  in  the  atmosphere  are  approximately  :  — 

Helium,  3  to  4  parts  per  million, 
Neon,  1  to  2  parts  per  hundred  thousand, 
Krypton,  1  part  in  20,000,000, 
Xenon,  1  part  in  170,000,000. 

Helium  was  detected  in  the  atmosphere  of  the  sun  by 
Lockyer  in  1868.  It  was  found  by  Ramsay,  soon  after  he 
discovered  argon,  in  the  gases  expelled  from  certain  minerals 
and  in  the  gas  and  water  of  certain  mineral  springs.  Helium 
is  one  of  the  disintegration  products  of  radium.  Niton  (Nt) 
belongs  to  this  group  of  elements  (see  Radioactivity). 

The  process  of  separating  the  inert  gases  of  the  atmosphere 
from  each  other  is  an  excellent  illustration  of  modern  ex- 


124  INORGANIC  CHEMISTRY 

perimental  methods.  The  mixture  of  the  five  gases  is 
compressed  in  a  bulb  and  cooled  to  about  —  185°  C.  by 
immersion  in  liquid  air;  the  argon,  krypton,  and  xenon 
condense  to  a  liquid  in  which  the  helium  and  neon  dissolve. 
When  the  bulb  is  removed  and  warmed,  the  helium  and  neon 
together  with  considerable  argon  escape  first  into  a  special 
bulb,  the  argon  next,  and  finally  the  krypton  and  xenon. 
Several  repetitions,  however,  are  necessary  to  separate  the 
argon  from  the  helium  and  neon  and  the  krypton  and  xenon, 
as  well  as  the  last  two  from  one  another.  By  immersing  the 
bulb  containing  the  helium  and  neon  in  liquid  hydrogen, 
the  neon  solidifies  and  the  helium  can  be  removed  first  by 
a  pump;  subsequently  the  neon  when  warmed  can  be  simi- 
larly removed  as  the  pure  gas. 

When  these  inert  gases  are  examined  by  a  spectroscope, 
they  exhibit  striking  spectra.  That  is,  when  electric  sparks 
are  passed  through  a  closed  tube  containing  any  of  these 
gases  and  the  light  thereby  produced  is  viewed  through  a 
spectroscope,  many  colored  vertical  lines  are  seen.  Certain 
lines  are  conspicuous,  e.g.  the  orange  line  in  the  case  of 
helium,  and  by  means  of  these  and  other  lines  it  is  possible 
to  detect  small  quantities  of  these  gases  and  to  distinguish 
them  from  other  gases,  since  no  two  spectra  are  exactly 
alike  under  given  conditions.  (For  an  account  of  the  spec- 
troscope and  its  application,  see  Chapter  XXV.) 

Air  is  a  mixture,  in  spite  of  the  fact  that  we  speak  of  its 
"  composition. "  Chemical  compounds,  as  we  have  already 
learned,  have  two  invariable  characteristics;  viz.  (1)  their 
constituents  are  in  a  fixed  proportion,  and  (2)  their  formation 
and  decomposition  are  usually  attended  by  definite  evidences 
of  chemical  action,  such  as  light,  heat,  electrical  phenomena, 
change  of  color,  etc.  The  following  facts  show  that  air  ig 
a  mixture  of  free  gases  ;  — 


THE   ATMOSPHERE  125 

(1)  The  proportion  of  oxygen  and  nitrogen  in  the  air  is 
not  fixed,  but  varies,  though  between  very  narrow  and  deter- 
minable  limits. 

(2)  When  nitrogen  and  oxygen  are  mixed  in  the  propor- 
tions which  form  air,  the  product  is  exactly  like  air,  but  the 
act  of  mixing  gives  no  evidence  of  chemical  action. 

(3)  When  air  is  dissolved  in  water,  a  greater  proportion 
of  oxygen   than   nitrogen   dissolves;  i.e.   they    dissolve   as 
independent  gases  in  proportions  fixed  by  their  intrinsic  solu- 
bility and  partial  pressure.    (See  Solubility  of  Gases.)     If  the 
oxygen  and  nitrogen  were  combined  in  the  air,  the  dissolved 
air  would,   of   course,   have   the   same  composition  as  air 
itself. 

(4)  Oxygen    and     nitrogen     distill     independently   from 
liquid    air. 

Liquid  air  is  a  mixture  of  the  liquefied  gases  which  con- 
stituted the  air  used.  It  is  a  milky  liquid,  owing  to  the 
presence  of  solid  carbon  dioxide  and  ice.  If  these  solids 
are  removed  by  filtering,  the  filtrate  has  a  pale  blue  tint. 
It  is  intensely  cold,  and  boils  at  about  —  190°  C.  under 
atmospheric  pressure.  If  a  vessel  is  filled  with  liquid  air, 
the  latter  boils  vigorously,  the  surrounding  air  becomes 
very  cold,  frost  gathers  on  the  vessel,  and  in  a  short  time 
the  liquid  air  will  have  entirely  disappeared  into  the  air  of 
the  room.  If,  however,  the  liquid  air  is  placed  in  a  Dewar 
bulb,  the  evaporation  is  only  slightly  affected  by  changes 
of  temperature. 

The  Dewar  bulb  (Fig.  14)  consists  of  two  flasks,  one 
within  the  other,  sealed  together  by  an  air-tight  joint  at  the 
top;  the  space  between  the  flasks  is  a  vacuum.  Sometimes 
the  surfaces  of  the  flasks  are  coated  with  silver,  which  re- 
flects the  heat  and  thereby  retards  the  evaporation  of 
the  contents. 


126 


INORGANIC   CHEMISTRY 


Liquid  air,  owing  to  its  extremely  low  temperature, 
produces  remarkable  physical  changes.  A  tin  or  iron  vessel 
which  has  been  cooled  by  liquid  air  is  so  brittle  that  it  may 
often  be  crushed  with  the  fingers.  Nearly  all  plastic  or 
soft  substances,  including  many  kinds  of  food,  when  im- 
mersed in  liquid  air  become  hard  and  brittle,  leather  being 


V 


Fia.  14.  —  Dewar  bulbs. 

the  only  important  exception.  Mercury  freezes  so  hard 
in  liquid  air  that  it  can  be  used  as  a  hammer  to  drive  a  nail. 
When  liquid  air  is  put  in  a  teakettle  standing  on  a  block 
of  ice  the  liquid  air  boils  vigorously.  If  the  kettle  of  liquid 
air  is  placed  over  a  lighted  Bunsen  burner,  frost  and  ice 
collect  on  the  bottom  of  the  kettle,  because  the  intense  cold 
of  the  kettle  solidifies  the  water  vapor  and  carbon  dioxide 
which  are  the  two  main  products  of  burning  illuminating 
gas.  If  water  is  now  poured  into  the  kettle,  the  liquid 
air  boils  vigorously  and  the  water  is  quickly  frozen;  the 
water  is  so  much  hotter  than  the  liquid  air  that  the  latter 
boils  more  violently,  and  since  its  rapid  evaporation  causes 
absorption  of  heat,  the  water  loses  its  heat  and  becomes 
ice.  Ordinary  liquid  air  is  from  one  fifth  to  one  half  liquid 
oxygen,  and  will  support  combustion.  A  red-hot  rod  of 
steel  or  of  carbon  burns  brilliantly  in  this  cold  liquid. 


THE   ATMOSPHERE  127 

Numerous  applications  of  liquid  air  have  been  proposed, 
but  thus  far  they  have  not  passed  the  experimental  stage. 
It  has  been  suggested  that  it  be  used  as  a  refrigerant  instead 
of  ice,  for  ventilating  and  cooling  rooms,  as  a  blasting 
material,  for  removing  diseased  flesh  from  a  wound,  for  de- 
stroying refuse,  and  as  a  source  of  oxygen  and  nitrogen.  The 
last  use  is  based  primarily  on  the  fact  that  as  liquid  air 
evaporates  the  nitrogen  passes  off  first,  and  in  a  short  time 
relatively  pure  oxygen  remains.  (See  Oxygen.) 

A  little  liquid  air  was  produced  in  1883  with  considerable 
labor  and  at  an  enormous  expense.  At  present  it  is  easily 
manufactured  in  large  quantities  at  a  comparatively  low 
cost.  Compressed  air  cooled  by  water  is  forced  through 
a  pipe  to  a  valve.  As  it  escapes  through  the  valve,  it  expands 
and  its  temperature  falls,  because  expansion  is  a  cooling 
process.  After  expansion,  the  cold  air  is  led  back  over  the 
outer  surface  of  the  same  pipe  by  which  it  came,  whereupon 
it  rapidly  regains  its  former  temperature.  But  in  doing 
so  it  cools  the  pipe  itself  and  the  air  within  it.  This  latter 
air  in  turn  expands  and  falls  in  temperature,  but  as  it  was 
cooler  than  the  first  portion  before  expansion,  so  it  is  colder 
than  it  after  expansion.  As  the  pressure  within  the  pipe 
is  maintained  by  a  continuous  supply  of  compressed  air, 
the  pipe  becomes  continually  colder  until  finally  the  expand- 
ing air  at  the  valve  liquefies  in  part  and  is  collected  in  a 
suitable  receptacle. 

Liquefaction  of  Gases.  —  In  the  first  methods  used  to 
liquefy  gases,  the  gas  was  subjected  simultaneously  to  a 
great  pressure  and  a  low  temperature.  Thus,  Faraday 
about  1823  liquefied  chlorine  gas  by  heating  one  arm  of  a 
sealed  bent  tube  containing  a  chlorine  compound  after 
having  placed  the  other  end  in  a  Breezing  mixture  ;  the 
liberated  chlorine  being  unable  to  escape  was  liquefied  by 


128  INORGANIC  CHEMISTRY 

the  pressure  and  low  temperature.  Other  gases  were 
liquefied  by  a  similar  method.  A  few,  however,  could  not 
be  liquefied,  e.g.  oxygen,  hydrogen,  and  nitrogen,  and  these 
were  called  permanent  gases.  About  1870  it  was  shown 
that  if  these  so-called  permanent  gases  were  cooled  to  a 
sufficiently  low  temperature,  they  could  be  liquefied  if  the 
pressure  was  also  sufficiently  great.  It  was  found,  further- 
more, that  each  gas  has  a  critical  temperature,  i.e.  a  tem- 
perature below  which  it  must  be  cooled  to  produce  lique- 
faction, no  matter  what  the  pressure.  Thus,  the  critical 
temperature  of  oxygen  is  about  — 118°  C.,  that  of  carbon 
dioxide  is  about  +  31°  C.,  that  of  normal  air  is  about  —  140° 
C.,  and  that  of  sulphur  dioxide  is  about  +  155.5°  C.  Ob- 
viously, critical  temperatures  vary  between  wide  limits. 
The  pressure  which  must  be  applied  to  liquefy  a  gas  at  its 
critical  temperature  is  called  its  critical  pressure.  The 
latter  value  varies,  but  not  between  such  wide  limits  as  the 
critical  temperature.  It  is  about  58  atmospheres  for  oxygen, 
15.3  for  hydrogen,  and  113  for  ammonia  (NH3).  As  the 
temperature  falls  below  the  critical  point,  less  pressure  than 
the  critical  amount  is  needed  for  liquefaction,  and  if  the 
temperature  of  the  gas  is  reduced  to  the  boiling  point  of  the 
liquefied  gas,  no  external  pressure  whatever  is  needed  for 
liquefaction.  Hence  the  essential  point  in  liquefying  most 
gases  is  the  production  of  a  sufficiently  low  temperature. 
This  can  be  done  in  some  cases  by  external  application  of 
cold,  though  in  the  case  of  gases  having  a  low  critical  tem- 
perature/the  cooling  is  now  produced  by  a  purely  mechanical 
process,  e.g.  like  that  used  for  liquefying  air  (see  above).  By 
this  process  all  known  gases  have  been  liquefied. 

If  a  liquefied  gas  can  be  cooled  to  a  sufficiently  low  tem- 
perature, it  becomes  solid.  Thus,  Dewar  by  boiling  liquid 
hydrogen  under  reduced  pressure  froze  it  to  a  foam-like 
solid,  the  temperature  being  about  —258°  C. 


THE  ATMOSPHERE  129 

NITROGEN 

Occurrence.  —  Nitrogen,  besides  comprising  four  fifths 
of  the  atmosphere,  is  a  constituent  of  nitric  acid  and  am- 
monia, and  of  many  compounds  related  to  them.  It  is 
also  an  essential  constituent  of  animal  and  vegetable  matter. 

The  name  nitrogen  was  given  to  the  gas  by  Chaptal  from  the  fact 
that  it  is  a  constituent  of  niter,  an  old  name  of  potassium  nitrate. 

Preparation.  —  Nitrogen  can  be  obtained  from  the  air 
by  removing  the  oxygen  by  phosphorus.  A  tall  jar  is 
placed  over  burning  phosphorus  contained  in  a  shallow 
dish  floating  in  a  large  vessel  of  water.  The  oxygen  com- 
bines with  the  phosphorus,  leaving  nitrogen,  more  or  less 
pure,  in  the  jar.  Other  methods  may  be  used,  such  as 
passing  air  over  heated  copper,  or  decomposing  ammonium 
dichromate  by  heat.  It  is  prepared  in  the  laboratory  by 
heating  a  mixture  of  sodium  nitrite  and  ammonium  chloride. 
It  can  be  prepared  commercially  from  liquid  air  (see  page  17). 

Additional  Properties.  —  In  addition  to  its  inertness, 
already  mentioned,  nitrogen  is  a  little  lighter  than  air, 
and  is  very  sparingly  soluble  in  water.  Its  density  is  .972 
(air  =  1).  One  liter  at  0°  C.  and  760  mm.  weighs  1.2507 
gm.  One  hundred  liters  of  water  dissolve  only  about  1.5  1. 
at  the  ordinary  temperature.  The  critical  temperature 
is  about  -146°  C.  Liquid  nitrogen  boils  at  -195.5°  C. 
under  ordinary  atmospheric  pressure,  and  solid  nitrogen 
melts  at  about  -214°  C. 

It  combines  directly  with  silicon  and  also  with  many  metals 
at  a  red  heat,  forming  nitrides,  e.g.  magnesium  nitride 
(Mg3N2) .  At  high  temperatures  and  under  special  conditions 
nitrogen  combines  with  oxygen  and  with  hydrogen,  forming 
nitric  oxide  (NO)  and  ammonia  (NH3).  (See  pages  212, 
217.) 


130 


INORGANIC  CHEMISTRY 


Relation  of  Nitrogen  to  Life.  —  Oxygen,  carbon  dioxide, 
and  water  vapor  are  essentially  related  to  the  life  of  plants 
and  animals.  Nitrogen  is  also  vitally  connected  with 
different  forms  of  life.  Atmospheric  nitrogen  merely  dilutes 
the  oxygen.  Although  we  live  in  an  atmosphere  contain- 
ing such  a  large  proportion  of  nitrogen,  we  cannot  assimilate 
it.  According  to  a  reliable  authority,  "  the  air  as  it  leaves 
the  lungs  contains  79.5  per  cent  of  nitrogen,"  and  hence 
cannot  become  a  part  of  the  body.  Yet  all  flesh  contains 
nitrogen,  and  certain  rejected  waste  products  of  animals 


A  B 

FIG.  15.  —  A  leguminous  plant  (the  hairy  vetch)  with  (B)  and  without 
(A)  nodules  on  the  roots. 

contain  considerable  combined  nitrogen.  The  nitrogen 
needed  by  animals  must  be  in  combination  to  become  avail- 
able. And  it  is  taken  in  the  form  of  nitrogenous  food 
such  as  lean  meat,  fish,  wheat  and  other  grains. 

Most  plants  take  up  combined  nitrogen  from  the  soil  in 
the  form  of  nitrates  (compounds  derived  from  nitric  acid) 
or  of  ammonia.  Hence  combined  nitrogen  is  being  con- 


THE  ATMOSPHERE  131 

stantly  taken  from  the  soil,  and  in  order  to  preserve  the  fer- 
tility of  the  soil,  nitrogen  must  be  supplied.  This  is  done 
by  adding  to  the  soil  a  fertilizer  containing  nitrogen  com- 
pounds. Sometimes  nitrogenous  organic  matter  is  used, 
such  as  manure,  dried  blood,  and  meat  or  fish  scraps.  Chem- 
ical fertilizers  are  extensively  used,  e.g.  sodium  nitrate 
(NaNO3),  ammonium  sulphate  ((NH4)2SO4),  calcium  nitrate 
(Ca(NO3)2),  or  calcium  cyanamide  (CaN2C).  Leguminous 
plants,  such  as  peas,  beans,  and  clover,  assimilate  free  nitro- 
gen directly  from  the  air  by  means  of  bacteria  which  are  in 
nodules  on  their  roots  (Fig.  15).  This  process  is  called  fixa- 
tion of  nitrogen.  Sometimes  soils  are  treated  with  a  prep- 
aration which  contains  "  nitrogen  bacteria." 

PROBLEMS  AND  EXERCISES 

1.  A  quantity  of  air  measures  24  1.  at  15°  C. ;  the  temperature 
is  reduced  to  —  16°  C.     What  is  now  the  volume?     (Pressure  un- 
changed.) 

2.  If  100,000  cu.  m.  of  air  at  any  pressure  were  saturated  with 
moisture  at  20°  C.,  what  weight  of  water  is  deposited  at  0°  C.? 

3.  If  air,  at  760  mm.,  has  a  temperature  of  20°  C.,  and  its  dew- 
point,  i.e.  the  temperature  at  which  it  is  saturated  with  moisture, 
is  15°  C.,  what  per  cent  of  moisture  by  volume  does  it  contain? 

4.  If  1  cc.  of  dry  air,  under  standard  conditions,  weighs  .00129 
gm.,  what  would  be  the  weight  of  1  cc.  of  air  saturated  with  mois- 
ture at  30°  C.  under  normal  pressure? 

5.  How  many  (a)  grams  and  (6)  cubic  centimeters  of  argon  can 
be  obtained  from    1500    kg.    of   pure   dry   air?     (Standard   con- 
ditions.) 

6.  How  many  grams  of  copper  oxide  (CuO)  are  formed  by  pass- 
ing 1728  gm.  of  normal  air  over  pure  copper?     How  many  grams 
of  magnesium  nitride  (MgsNa)  by  passing  the  residual  gas  over  red- 
hot  magnesium?     How  many  grams  of  gas  remain?     Equations 
are  Cu  +  O  =  CuO  and  3  Mg  +  2  N  =  Mg3N2. 

7.  Write  equations  for  the  following  reactions  :    (a)  Phosphorus 
pentoxide  and  water  form  phosphoric  acid  (H3PO4).     (b)  Copper 
oxide  (CuO)  and  hydrogen  from  copper  and  water. 


CHAPTER  IX 
Solution  —  Theory  of  Electrolytic  Dissociation 

Introduction.  —  Many  facts  about  solutions  have  already 
been  stated  (see  Chapter  V).  The  present  chapter  is  a  con- 
tinuation of  the  discussion  of  solutions  with  special  reference 
to  the  theory  of  their  nature  and  the  interpretation  of  certain 
phenomena  by  this  theory.  The  term  solution  will  be  re- 
stricted to  aqueous  solutions,  i.e.  those  in  which  water  is  the 
solvent. 

General  Properties  of  Solutions.  — It  is  very  desirable  to 
recall  at  this  point  certain  properties  of  solutions.  First, 
the  solubility  of  the  solute  in  the  solvent  varies  widely  and  is 
generally  increased  by  rise  of  temperature  until  it  reaches  a 
limit,  thereby  giving  the  phenomena  connected  with  unsatu- 
ration,  saturation,  and  supersaturation.  Second,  the  solute 
in  most  cases  can  be  recovered  unchanged  by  evaporating, 
cooling,  or  distilling  the  solution.  Third,  the  solution  often 
contains  the  solute  in  an  especially  favorable  condition  for 
chemical  action.  Finally,  solutions  have  a  definite  boiling 
point,  freezing  point,  and  vapor  pressure,  which  differ  from 
the  corresponding  values  of  the  solvent. 

Behavior  of  Solutions  toward  an  Electric  Current.  —  Pure 
water  and  pure  dry,  solid  sodium  chloride  do  not  appreciably 
permit  the  passage  of  an  electric  current.  But  a  solution  of 
sodium  chloride  is  an  excellent  conductor  of  electricity,  and 
the  same  is  true  of  a  solution  of  hydrochloric  acid  and  of 
sodium  hydroxide.  On  the  other  hand,  sugar  is  a  non-con- 
ductor, both  in  solution  and  in  the  solid  state.  In  a  word, 

132 


SOLUTION  133 

solutions  of  certain  substances  conduct  electricity,  while 
solutions  of  others  do  not.  Hence  solutions  can  be  divided 
on  this  basis  into  two  classes,  viz. :  (1)  electrolytic  solutions, 
or  those  which  conduct  electricity,  and  (2)  non-electrolytic 
solutions,  or  those  which  do  not  conduct  electricity.  Sub- 
stances whose  solutions  are  electrolytic  are  called  electro- 
lytes, and  those  whose  solutions  are  not  electrolytic  are  called 
non-electrolytes.  Each  class  has  characteristic  properties. 
But  electrolytes  possess  certain  conspicuous  properties 
which  are  not  exhibited  by  non-electrolytes,  and  it  was  the 
study  of  these  characteristic  properties  that  led  to  the  pro- 
posal of  the  present  theory  of  solutions.  It  is  called  the 
theory  of  electrolytic  dissociation  and  was  proposed  by  the 
Swedish  physicist  Arrhenius  in  1887;  its  general  adoption 
has  been  hastened  by  the  work  of  Van't  Hoff,  Ostwald,  and 
Nernst. 

Theory  of  Electrolytic  Dissociation.  — It  was  believed  for 
many  years  that  the  molecules  of  a  dissolved  substance  were 
distributed  unchanged  throughout  the  solvent.  It  was 
also  believed  that  the  molecules  of  certain  dissolved  sub- 
stances combined  to  some  degree  with  the  molecules  of  the 
solvent.  Evidence  is  fast  accumulating  which  indicates  that 
in  many  solutions  the  solute  is  in  the  form  of  molecules, 
in  others  the  solvent  does  unite  with  the  solute  or  some 
>f  its  constituents.  The  present  theory  of  solutions  differs 
from  these  by  offering  an  explanation  of  solution  which  is 
>re  comprehensive.  Briefly,  the  theory  of  electrolytic  dis- 

>ciation  assumes  (1)  that  molecules  of  electrolytes  when 
dissolved  in  water  break  up  to  a  varying  degree  into  inde- 
pendent particles  charged  with  electricity,  and  (2)  that  the 
nature  and  number  of  these  electrically  charged  particles 
determine  to  a  large  extent  certain  physical  and  chemical 
properties  of  solutions. 


134  INORGANIC  CHEMISTRY 

Before  stating  the  facts  on  which  the  theory  is  based,  it 
will  be  necessary  to  expand  the  two  assumptions  and  to  de- 
fine several  terms.  The  breaking  up  of  certain  substances 
when  in  aqueous  solution  is  called  electrolytic  dissociation, 
or  ionization.  The  independent  particles  are  called  ions. 
Thus,  when  sodium  chloride  is  dissolved  in  water,  some  of  its 
molecules  dissociate  into  sodium  ions  and  chlorine  ions. 
Each  ion  is  a  portion  of  a  molecule  and  is  charged  with  elec- 
tricity. Two  kinds  of  ions  are  present  in  every  electrolytic 
solution,  viz.  electro-positive  ions,  or  cations,  and  electro- 
negative ions,  or  anions.  Ions,  although  formed  by  the 
dissociation  of  molecules,  are  not  identical  with  atoms,  but 
differ  mainly  in  having  a  charge  of  electricity.  For  exam- 
ple, when  sodium  chloride  is  dissolved  in  water,  the  electro- 
positive sodium  ions  move  about  in  the  water  without  pro- 
ducing any  apparent  chemical  change;  but  ordinary  sodium 
interacts  violently  with  water,  as  we  have  already  seen. 
Similarly,  the  chlorine  ions  circulate  freely  in  water  and  ex- 
hibit none  of  the  effects  of  gaseous  chlorine  upon  water.  In 
a  word,  in  such  a  solution  the  sodium  ions  and  chlorine  ions 
exist  side  by  side  without  any  apparent  decomposition  of  the 
water  or  any  apparent  tendency  to  combine  with  each  other. 
The  properties  of  ions,  as  already  stated,  are  mainly  due  to 
their  electric  charges,  and  ions  may  be  defined  as  electrically 
charged  atoms  or  atomic  groups.  It  is  customary  to  repre- 
sent ions  by  chemical  symbols  supplemented  by  the  sign 
which  designates  the  kind  of  electric  charge.  Thus,  the 
ions  formed  by  the  dissociation  of  sodium  chloride  are  Na+ 
and  Cl~,  while  potassium  nitrate  yields  K+  and  NO8~.  A 
solution  of  sodium  chloride  gives  no  evidence  of  electricity. 
In  general,  solutions  of  electrolytes  are  electrically  neutral; 
i.e.  although  they  allow  an  electric  current  to  pass  through 
them  when  supplied  from  some  external  source  such  as  a 
battery  or  dynamo,  the  solution  itself  is  electrically  neutral. 


SOLUTION  135 

It  therefore  follows  that  the  sum  of  the  electric  charges  of 
the  positive  ions  equals  the  sum  of  the  electric  charges  of 
the  negative  ions.  The  ions  balance  each  other  electrically. 
Thus,  in  the  case  of  sodium  chloride  solution,  the  number 
of  sodium  ions  equals  the  number  of  chlorine  ions  and  the 
sum  of  the  positive  charges  on  the  sodium  ions  equals 
the  sum  of  the  negative  charges  on  the  chlorine  ions.  In  the 
case  of  calcium  chloride  (CaCl2),  each  molecule  dissociates 
into  two  chlorine  ions  and  one  calcium  ion;  but  since  the 
sum  of  each  kind  of  electric  charges  must  be  the  same,  each 
calcium  ion  must  carry  twice  the  charge  which  is  on  each  chlo- 
rine ion.  Hence  the  calcium  ions  are  designated  as  Ca++  and 
the  chlorine  ions  as  2  Cl~.  A  unit  charge  of  electricity  is 
indicated  by  the  single  sign,  and  multiples  by  the  proper 
number.  For  example,  the  sodium  ion  is  Na+,  the  common 
copper  ion  is  Cu++,  the  aluminium  ion  is  Al+++,  the 
sulphate  ion  is  SO4  ,  and  the  nitrate  ion  is  NO3".  The 
degree  of  dissociation  varies  with  the  concentration  and  with 
the  electrolyte.  The  greater  the  dilution,  the  greater  the 
dissociation.  Conversely,  dissociation  is  decreased  by  de- 
creasing the  volume  of  the  solvent;  the  ions  tend  to  unite, 
forming  undissociated  molecules  which  can  be  ultimately 
obtained  as  a  mass  of  the  original  substance  by  evaporation. 
Consequently,  a  solution  contains  undissociated  molecules 
as  well  as  ions,  depending  upon  the  substance  and  the  con- 
centration of  the  solution.  Experiment  shows  that  the  chem- 
ical behavior  of  a  dissolved  substance  often  depends  largely 
on  the  extent  of  the  dissociation  of  the  molecules  into  ions. 

As  previously  stated,  only  certain  substances  are  electro- 
lytes. These  are  acids,  bases,  and  salts.  The  general  prop- 
erties of  these  substances  are  discussed  in  Chapter  X  and 
their  special  characteristics  are  treated  under  the  individual 
compounds.  It  is  sufficient  for  our  present  purpose  to  re- 
gard them  as  a  single  class  of  substances  which  dissociate 


136  INORGANIC  CHEMISTRY 

in  solution  into  ions.  That  is,  ions  are  in  the  solutions  of  the 
familiar  acids  like  sulphuric,  hydrochloric,  and  nitric,  the 
familiar  bases  like  sodium  hydroxide  and  potassium  hydrox- 
ide, and  the  familiar  salts  like  potassium  chlorate,  sodium 
chloride,  sodium  sulphate,  and  many  others  which  will  soon 
be  described.  Sugar,  alcohol,  and  other  compounds  not  so 
familiar  do  not  dissociate  into  ions  when  dissolved  in  water 
and  their  solutions  do  not  conduct  electricity.  Neverthe- 
less, such  solutions  have  specific  and  instructive  properties, 
especially  when  compared  with  the  properties  of  electro- 
lytic solutions. 

Properties  of  Electrolytes  and  Non-electrolytes.  — The 
theory  of  electrolytic  dissociation,  which  has  just  been  dis- 
cussed, is  based  on  many  facts  which  have  accumulated  as 
the  outcome  of  a  comprehensive  study  of  the  properties  of 
solutions,  both  electrolytic  and  non-electrolytic.  These  facts 
will  now  be  presented.  Careful  distinction  should  be  drawn 
between  the  statements  in  the  paragraphs  immediately  pre- 
ceding and  those  about  to  be  made.  The  preceding  con- 
cerned the  theory  of  electrolytic  dissociation.  The  forth- 
coming concern  facts  and  laws  derived  from  experiment  and 
interpreted  by  the  theory. 

(a)  Osmotic  Pressure.  —  The  passage  of  a  liquid  through  a 
membrane  is  called  osmosis.  Certain  membranes  permit 
the  passage  of  the  solvent  but  prevent  more  or  less  the  pas- 
sage of  the  solute ;  such  membranes  are  said  to  be  semiper- 
meable.  Osmosis  is  a  common  phenomenon  in  physiological 
processes,  for  semipermeable  membranes  occur  in  animal  and 
vegetable  organisms.  Osmosis  and  osmotic  pressure  can  be 
demonstrated  by  a  simple  experiment.  A  piece  of  parchment 
paper  is  tied  tightly  over  the  larger  end  of  a  thistle  tube, 
which  is  then  partly  filled  with  a  concentrated  sugar  solution 
and  immersed  in  a  vessel  of  water  (Fig.  15  a) .  The  membrane 


SOLUTION 


137 


is  permeable  to  water,  but  not  to  sugar.  Water  passes 
through  the  membrane  into  the  sugar  'solution,  which  in- 
creases in  volume  and  hence  rises  in  the  tube.  If  the  mem- 
brane is  strong  enough,  the  column  of  solution  will  rise  to  a 
maximum  height.  At  this  point  the 
weight  of  the  solution,  and  therefore 
its  pressure,  is  such  that  the  tendency 
of  the  water  to  pass  through  the  mem- 
brane into  the  solution  equals  the  ten- 
dency of  the  water  to  pass  out.  When 
this  condition  is  reached,  the  weight  of 
the  liquid  above  the  membrane  is  a 
measure  of  the  osmotic  pressure  of  the 
sugar  solution. 

In  accurate  measurements  of  the 
osmotic  pressure  of  aqueous  solutions  the 
semipermeable  membrane  is  usually  a 
film  of  cupric  ferrocyanide  (Cu2Fe(CN)6)  FlG- is  a.  —  Experiment 
deposited  in  the  pores  of  an  unglazed 
porcelain  vessel.  The  cell,  as  it  is  called, 
is  filled  with  the  solution  to  be  investigated,  fitted  tightly 
with  a  special  form  of  stopper,  connected  with  a  manometer 
(to  measure  the  pressure),  and  then  immersed  in  water. 
Water  flows  slowly  through  the  membrane  into  the  cell 
until  a  maximum  pressure  is  produced ;  the  system  is  then  in 
equilibrium,  i.e.  the  tendency  of  the  water  to  pass  into  the 
solution  through  the  semipermeable  membrane  is  balanced 
by  the  opposing  pressure  of  the  manometer.  This  increase 
over  the  original  hydrostatic  pressure  is  read  on  the  manom- 
eter and  is  equal  to  the  osmotic  pressure  of  the  sugar  solu- 
tion. Determinations  of  the  osmotic  pressure  of  dilute  solu- 
tions of  non-electrolytes  show  that  osmotic  pressure  is 
directly  proportional  (1)  to  the  concentration  and  (2)  to  the 
absolute  temperature  of  the  solution ;  in  (1)  the  tempera- 


to   illustrate   osmotic 
pressure. 


138  INORGANIC  CHEMISTRY 

ture  must  be  kept  constant  and  in  (2)  the  volume  must  be 
kept  constant.  Furthermore,  if  a  gram-molecular  weight 
(that  is,  the  number  of  grams  numerically  equal  to  the  mo- 
lecular weight)  of  non-electrolytes  is  dissolved  in  equal  quan- 
tities of  water,  the  osmotic  pressure  of  each  solution  is  the 
same.  The  gram-molecular  weight  of  a  substance  is  called  a 
mole  —  a  convenient  value  which  is  frequently  used  in  stating 
concentration  and  in  comparing  experimental  results.  In  the 
case  of  osmotic  pressure,  for  example,  one  mole  (342  gm.)  of 
cane  sugar  (C^H^On)  and  one  mole  (58  gm.)  of  acetone 
(CaHeO)  dissolved  in  equal  quantities  of  water  have  the  same 
osmotic  pressure. 

The  osmotic  pressure  of  an  electrolytic  solution  is  found 
to  be  greater  than  that  of  a  non-electrolytic  solution 
under  the  same  conditions.  The  excess  of  pressure  varies 
somewhat  with  the  conditions,  i.e.  with  concentration,  tem- 
perature, and  the  substance.  Nevertheless,  the  difference 
between  the  normal  and  the  abnormal  osmotic  pressure  is 
conspicuously  large  in  the  case  of  many  substances  yielding 
electrolytic  solutions,  these  substances  being  acids,  bases, 
and  salts.  Thus,  a  solution  of  potassium  chloride  has  an 
osmotic  pressure  1.88  times  that  of  a  corresponding  sugar 
solution,  while  a  calcium  chloride  solution  has  a  pressure 
about  2.5  times.  This  discrepancy  between  the  values  of  the 
osmotic  pressure  of  non-electrolytic  and  electrolytic  solutions 
is  readily  explained,  if  the  facts  are  interpreted  by  the  theory 
of  electrolytic  dissociation.  According  to  this  theory  a  non- 
electrolytic  solution  contains  only  molecules  and  an  elec- 
trolytic solution  contains  both  ions  and  molecules.  Now 
osmotic  pressure  is  believed  to  be  due  to  the  independent 
particles  of  the  solute  in  solution,  and  the  amount  of  pressure 
is  determined  by  the  number  of  particles  in  a  given  volume. 
Hence,  when  equivalent  solutions  are  used,  the  electrolytic 
solution  contains  more  particles  than  the  non-electrolytic 


SOLUTION 


139 


because  some  of  the  molecules  in  the  electrolytic  solution 
are  dissociated  into  ions.  Facts  and  theory  agree  as  far  as 
osmotic  pressure  is  concerned. 

(b)  Freezing  Point  and  Boiling  Point  of  Solutions.  —  In 
the  preliminary  discussion  of  the  properties  of  solutions  given 
in  Chapter  V,  it  was  stated  that  the  freezing  point  of  a  solu- 
tion is  lower  than  the  freezing  point  of  the  solvent  and  that 
the  depression  of  the  freezing  point  is  proportional  to  the 
concentration  of  the  solution.  These  facts  can  be  expressed 
by  a  concrete  case,  thus :  — 

DEPRESSION  or  THE  FREEZING  POINT 


SOLVENT 

SOLUTE 

FREEZING  POINT 

DEPRESSION 

100  gm. 

water 
in  each  case 

11.4  gm.  sugar 
22.8  gm. 
34.2  gm. 

-  .62°  C. 
-1.23°C. 
-1.85°  C. 

.62 
2  x  .62  (approx.) 
3  x  .62  (approx.) 

Furthermore,  experiment  shows  that  the  freezing  point  of 
water  is  depressed  the  same  number  of  degrees,  if  1000  gm.  of 
water  contain  a  mole  of  certain  solutes  (i.e.  a  number  of 
grams  numerically  equal  to  the  molecular  weight).  Thus, 
the  freezing  point  of  water  is  depressed  about  1.86°  C.  by  a 
solution  of  342  gm.  of  cane  sugar  (C^H^On)  and  58  gm.  of 
acetone  (C3H6O),  each  dissolved  in  1000  gm.  of  water.  This 
uniform  behavior  is  not  well  exhibited,  however,  unless  the 
solutions  are  dilute  and  involve  no  chemical  action  between 
solvent  and  solute.  On  the  other  hand,  when  solutions  of 
electrolytes,  i.e.  acids,  bases,  and  salts,  are  experimented 
with,  the  freezing  point  is  lower  than  that  produced  by  non- 
electrolytes  under  the  same  experimental  conditions.  More- 
over, the  depression  is  not  uniform  for  all  electrolytic  solu- 
tions under  uniform  conditions.  For  example,  a  solution 
containing  a  mole  of  sodium  chloride  (i.e.  58.5  gm.)  in  1000 


140  INORGANIC  CHEMISTRY 

grams  depresses  the  freezing  point  of  water  about  3.5°  C.,  or 
nearly  twice  the  amount  produced  by  a  cane  sugar  solution 
of  equivalent  concentration.  This  exceptional  behavior  of 
solutions  of  acids,  bases,  and  salts  can  be  explained  as  in  the 
case  of  osmotic  pressure.  Solutions  of  non-electrolytes  con- 
tain only  molecules,  while  solutions  of  electrolytes  contain 
ions  into  which  some  of  the  molecules  have  dissociated. 
Hence  the  number  of  independent  particles  (molecules  and 
ions)  in  the  electrolytic  solution  is  greater  than  in  the  non- 
electrolytic  solution.  Ions  and  molecules  act  alike  on  the 
freezing  point  of  a  solution,  and  the  larger  the  number  of 
particles,  the  greater  the  depression.  This  deduction  is 
further  confirmed  by  the  fact  that  electrolytic  solutions  in 
which  a  large  proportion  of  molecules  is  dissociated  show  a 
relatively  greater  depression  than  those  in  which  the  disso- 
ciation is  limited  —  the  degree  of  dissociation  being  found  by 
an  independent  experiment. 

Analogous  statements  can  be  made  about  the  elevation  of 
the  boiling  point  of  solutions.  That  is,  (1)  the  boiling  point 
of  a  solution  is  higher  than  that  of  the  solvent,  (2)  the  ele- 
vation of  the  boiling  point  is  proportional  to  the  concentra- 
tion, and  (3)  the  elevation  is  the  same  (i.e.  .52°  C.)  in  the  case 
of  all  non-electrolytic  solutions  containing  a  mole  of  the  solute 
in  1000  gm.  of  water.  But  solutions  of  acids,  bases;  and  salts 
behave  exceptionally.  They  boil  at  a  higher  temperature 
than  non-electrolytic  solutions  of  the  same  concentration. 
The  explanation  offered  by  the  theory  of  electrolytic  disso- 
ciation is  the  same  as  in  the  preceding  cases,  viz.  dissociation 
of  some  of  the  molecules  into  ions,  which  affect  the  boiling 
point  in  the  same  way  as  the  molecules  themselves. 

(c)  Electrolysis  of  Solutions.  —  Electrolysis  is  the  series 
of  changes  accompanying  the  passage  of  an  electric  current 
through  a  solution.  It  is  accomplished  in  an  electrolytic  cell. 
This  piece  of  apparatus  consists  of  three  essential  parts  — 


SOLUTION 


141 


the   electrolytic   solution,   the   containing  vessel,   and   two 
electrodes  which  convey  the  electric  current  to  and  from  the 
solution.     A  simple  form  of  such  a  cell  is  shown  in  Figure  16. 
It  is   customary  to  speak  of  the 
current   as  entering  the   solution 
by  the  anode  or  positive  electrode 
and  as  leaving  by  the  cathode  or 
negative  electrode.     Both  the  elec- 
trolytic cell  and  the  ordinary  vol- 
cell    are    fully   described 


taic 


m 


convey  the  current  to  and 
from  the  electrolytic  solution 
(E) ;  B  or  D  is  the  battery 
or  dynamo  which  provides 
the  electricity. 


Chapter  XI,  though  a  general  idea  FIG.  16.  —  Simple  electrolytic 
of  the  electrolytic  cell  serves  our  cf;  A  (^nodf  *nd/  (c^h~ 

ode)  are  the  electrodes  which 

present  need. 

When  a  concentrated  solution  of 
hydrochloric  acid  gas  (commonly 
called  simply  hydrochloric  acid)  is 
put  into  an  electrolytic  cell  and 

subjected  to  the  action  of  an  electric  current,  two  gases 
are  liberated,  —  hydrogen  at  the  cathode  and  chlorine  at 
the  anode.  This  is  a  very  simple  illustration  of  electrolysis. 
Let  us  interpret  it  by  the  theory  of  electrolytic  dissociation. 
When  hydrochloric  acid  gas  is  dissolved  in  water,  hydrogen 
ions  (H+)  and  chlorine  ions  (Cl~)  are  immediately  formed 
by  the  dissociation  of  some  of  the  molecules  of  hydrochloric 
acid  (HC1).  These  ions  are  formed  in  the  solution  as  soon 
as  the  acid  dissolves  and  before  the  electric  current  is  con- 
nected with  the  cell.  As  soon  as  the  current  is  turned  on, 
however,  the  electrodes  at  once  become  charged  with  elec- 
tricity—  the  anode  assuming  a  positive  charge  and  the 
cathode  a  negative  charge.  Now  according  to  a  principle 
established  many  years  ago,  bodies  charged  with  like  kinds 
of  electricity  repel  each  other  and  bodies  charged  with  unlike 
kinds  attract  each  other.  Consequently  the  anions  or  electro- 
negative ions  move  toward  the  anode  or  electro-positive 


142  INORGANIC   CHEMISTRY 

electrode,  while  the  cations  or  electro-positive  ions  move 
toward  the  cathode  or  electro-negative  electrode,  or  briefly, 
"anions  to  anode,  cations  to  cathode.''  This  migration  of 
the  ions,  as  it  is  called,  toward  their  respective  poles  is  shown 
diagrammatically  in  Figure  17.  As  soon  as  the  ions  reach 

their  electrodes,  they  act  in 
accordance  with  another  long- 
established  principle;  that  is, 
they  give  up  their  electric 
charges.  In  other  words,  when 
the  electro-positive  cations  of 
_  _  hydrogen  touch  the  electro- 

FIG.  17.  —  Migration  of  ions  in  an  . 

electrolytic  cell.   The  cations  are  negative  cathode,  the  electric 
marked  H+  and  the  anions  ci~.  charges  are  neutralized.    Elec- 

A  is  the  anode  and  C  is  the  cathode.    tri(J  charges  are  constantly  re- 
B  or  D  is  the  battery  or  dynamo. 

newed  on  the  cathode  by  the 

battery  or  dynamo,  but  the  hydrogen  ions  once  deprived  of 
their  electric  charges  do  not  regain  them  and  immediately 
become  'ordinary,  uncharged  hydrogen  atoms,  which  combine 
and  escape  as  molecules  of  hydrogen  gas.  Similarly,  the  elec- 
tro-negative anions  of  chlorine  migrate  to  the  electro-positive 
anode,  lose  their  charges,  become  chlorine  atoms,  and  ulti- 
mately escape  as  chlorine  gas.  All  cases  of  electrolysis  are 
not  as  simple  as  this  one,  but  it  serves  admirably  as  an  intro- 
ductory illustration.  It  should  be  noted  that  the  electric 
current  does  not  tear  the  molecules  apart,  as  was  once  sup- 
posed. The  molecules  of  hydrochloric  acid  that  dissociate 
are  already  dissociated  before  the  electric  current  is  intro- 
duced. The  current  upsets  the  electrical  equilibrium  be- 
tween the  ions,  so  to  speak,  and  they  start  at  once  on  a  mi- 
gration toward  their  proper  electrodes  where  they  lose  their 
charges  and  become  ordinary,  uncharged  atoms  or  atomic 
groups.  Careful  and  extended  experiments  have  not  only 
demonstrated  the  actual  movement  of  ions,  but  have  deter- 


SOLUTION  143 

mined  the  rate  of  migration  in  many  cases.  Non-electrolytic 
solutions  do  not  conduct  electricity,  because,  according  to 
the  theory  of  electrolytic  dissociation,  they  contain  no  ions. 
In  this  connection  it  is  appropriate  to  emphasize  the  fact  that 
only  those  solutions  conduct  electricity  that  have  been  found 
by  other  methods  to  contain  ions. 

Electrolysis  is  often  a  complicated  process,  since  the  regen- 
erated atoms  and  atomic  groups  may  interact  chemically 
with  the  constituents  of  the  solution  and  sometimes  with 
each  other  or  with  the  electrodes.  The  electrolysis  of  copper 
sulphate  furnishes  a  typical  illustration.  The  ions  of  a 
copper  sulphate  solution  are  copper  ions  (Cu++)  and  sul- 
phate ions  (S04  ).  When  this  solution  is  electrolyzed, 
the  copper  ions  (Cu++)  lose  their  electric  charges  at  the 
cathode,  become  copper  atoms  (Cu),  and  adhere  as  metallic 
copper  to  the  cathode.  The  sulphate  ions  (SO4 )  lose 
their  electric  charges  at  the  anode  and  become  ordinary, 
uncharged  atomic  groups  (SO4) .  But  this  group  of  atoms  is 
chemically  unstable,  and  immediately  interacts  with  the 
water  around  the  anode,  forming  sulphuric  acid  (H2SO4)  and 
oxygen  (0).  The  oxygen  escapes,  but  the  sulphuric  acid 
mingles  with  the  solution  and  dissociates  into  its  ions. 
Similarly,  a  solution  of  sodium  sulphate  when  undergoing 
electrolysis  yields  sulphuric  acid  at  the  anode  and  sodium 
hydroxide  (NaOH)  at  the  cathode;  the  electrolyte  itself 
(sodium  sulphate)  furnishes  directly  only  sodium  and  sul- 
phate ions,  which  lose  their  charges  at  the  electrodes,  and 
by  their  subsequent  chemical  interaction  with  the  water  give 
the  final  result  just  stated. 

The  so-called  electrolysis  of  water  is  readily  interpreted 
by  the  theory  of  electrolytic  dissociation.  Water  itself  does 
not  conduct  electricity  to  an  extent  which  is  comparable  with 
the  behavior  of  an  electrolytic  solution,  because  water  dis- 
sociates only  inappreciably  and  gives  therefore  an  exceed- 


144  INORGANIC   CHEMISTRY 

ingly  small  number  of  ions.  A  solution  of  sulphuric  acid 
contains  hydrogen  ions  (H+H+)  and  sulphate  ions  (S04  ). 
When  a  current  is  passed  through  this  solution,  hydrogen 
ions  migrate  to  the  cathode,  lose  their  electric  charges,  be- 
come hydrogen  atoms,  and  eventually  escape  as  hydrogen 
gas;  the  SO4-ions  migrate  to  the  anode,  lose  their  electric 
charges,  become  SO4-groups,  and  interact  with  the  water 
to  form  sulphuric  acid  and  oxygen.  The  oxygen  escapes  as 
a  gas,  while  the  sulphuric  acid  dissociates  into  its  ions.  The 
water,  therefore,  is  not  split  up  directly  into  hydrogen  and 
oxygen,  as  was  formerly  supposed.  The  two  liberated 
gases  are  produced  by  the  joint  operations  of  electrolysis 
and  subsequent  chemical  action,  but  the  gases  would  not  be 
liberated  at  all  unless  the  ionization  of  the  sulphuric  acid 
had  previously  occurred  in  the  solution. 

Electrolysis  is  a  broad  subject,  and  is  not  limited  to  aqueous 
solutions.  In  subsequent  chapters  frequent  reference  will 
be  made  to  the  electrolysis  of  molten  substances,  especially 
to  the  industrial  applications  which  have  become  so  impor- 
tant. Enough  has  been  set  forth  at  present,  however,  to 
show  that  the  facts  thus  far  revealed  by  electrolysis  are  in 
harmony  with  the  theory  of  electrolytic  dissociation. 

(d)  Chemical  Behavior  of  Electrolytic  Solutions.  —  It  has 
been  pointed  out  that  chemical  action  is  often  dependent  upon 
the  presence  of  water.  Dry  compounds  like  potassium 
chloride  (KC1)  and  silver  nitrate  (AgNO3)  do  not  interact 
chemically,  but  if  their  solutions  are  mixed,  a  precipitate  of 
silver  chloride  (AgCl)  is  immediately  produced.  On  the 
other  hand,  there  is  no  chemical  action  manifested  when 
solutions  of  potassium  chlorate  (KC1O3)  and  silver  nitrate  are 
mixed,  despite  the  fact  that  chlorine  is  a  constituent  of 
potassium  chlorate.  Furthermore,  any  chloride  in  solution 
will  interact  with  silver  nitrate  in  solution  and  produce  a 
precipitate  of  silver  chloride.  These  facts  are  typical  of 


SOLUTION  145 

electrolytic  solutions.  Interpreted  by  the  theory  of  elec- 
trolytic dissociation,  they  mean  that  reactions  in  solutions 
are  due  to  some  extent  to  ions.  Dry  or  undissolved  electro- 
lytes do  not  interact,  because  no  ions  are  present,  but  in  the 
case  of  dissolved  electrolytes  certain  ions  at  once  seek  each 
other  out  in  accordance  with  the  fundamental  principles  of 
chemical  action.  If  this  action  results  in  the  formation  of 
an  insoluble  compound,  like  silver  chloride,  this  factor  is 
removed  from  the  scene  of  action  as  a  precipitate  and  serves 
as  visual  evidence  of  the  chemical  change.  Often  ions  are 
produced  which  cannot  enter  into  chemical  combination. 
Thus,  a  potassium  chlorate  solution  contains  potassium  ions 
(K+)  and  chlorate  ions  ((C103~),  and  when  silver  nitrate 
solution  is  added,  the  solution  contains  four  kinds  of  ions,  — 
potassium  ions  (K+),  chlorate  ions  (C103~),  silver  ions  (Ag+), 
and  nitrate  ions  (NO3~).  But  all  compounds  which  might 
be  formed  by  the  various  combinations  of  these  ions  are 
soluble.  Hence  the  ions  remain  as  such  in  the  solution.  It 
is  for  this  reason  that  silver  nitrate  is  effective  in  testing  for 
hydrochloric  acid  or  a  soluble  chloride,  but  not  for  other 
compounds  containing  chlorine,  such  as  potassium  chlorate 
(KC103)  and  chloroform  (CHC13).  Strictly  speaking,  the 
test  is  for  chlorine  ions  or  ionic  chlorine,  not  for  the  element 
chlorine;  and  since  the  solutions  of  potassium  chlorate  and 
chloroform  contain  no  ionic  chlorine,  the  test  fails  with 
these  compounds.  Similarly,  sulphuric  acid  and  all  soluble 
sulphates  form  insoluble  barium  sulphate  (BaSO4)  when 
added  to  a  solution  of  barium  chloride  (or  any  other  soluble 
barium  compound),  because  the  sulphuric  acid  and  sulphate 
solutions  contain  sulphate  ions  (S04~~),  which  combine 
with  the  barium  ions  (Ba++)  in  the  barium  chloride  solu- 
tion. But  other  sulphur  compounds,  such  as  sulphides, 
sulphites,  and  thiosulphates,  do  not  form  barium  sulphate 
when  added  to  barium  chloride  solution,  because  solutions 


146  INORGANIC  CHEMISTRY 

of  these  compounds  do  not  contain  sulphate  ions.  It  is  clear 
from  the  above  statements  why  a  single  test  (i.e.  the  pre- 
cipitation of  barium  sulphate)  is  applicable  to  sulphuric 
acid  and  all  soluble  sulphates.  All  contain  in  solution  a 
common  ion  (SO4  ). 

Other  properties  besides  the  formation  of  precipitates 
are  ascribed  to  ions  and  are  often  used  as  tests.  Thus,  the 
sour  taste  of  all  acids  is  attributed  to  hydrogen  ions  (H+), 
which  are  common  to  acids.  The  color  of  solutions  is  also 
due  to  ions.  Most  ions  are  colorless,  while  solutions  having 
a  common  colored  ion  have  the  same  color.  Thus,  copper 
ions  (Cu++)  are  blue,  and  solutions  containing  such  ions 
are  blue,  irrespective  of  the  color  of  the  undissolved  copper 
compound.  Cobalt  ions  (Co++)  are  pink,  and  nickel  ions 
(Ni++)  are  green  —  colors  usually  exhibited  by  solutions  of 
compounds  of  these  elements.  The  migration  of  ions  is  often 
studied  by  means  of  colored  ions. 

Common  Ions.  —  Ions,  as  already  stated,  are  electrically 
charged  atoms  or  atomic  groups.  It  is  rather  difficult  for 
a  beginner  to  determine  what  ions  are  present  in  a  solution. 
The  problem  is  simplified  somewhat  if  the  following  general 
statements  are  borne  in  mind,  (a)  Hydrogen  and  metals 
form  simple  cations,  (b)  Non-metals  (except  hydrogen)  form 
simple  anions.  (c)  Some  metals  (e.g.  Cr  and  Mn)  and  several 
non-metals  (e.g.  C,  N,  S,  P)  form  compound  ions  —  usually 
anions;  e.g.  HCO8-,  NO3~,  HSOr,  H2POr,  CrOr~,  MnOr 
(note  also  OH~  and  NH4+).  Certain  elements  likewise  form 
complex  ions;  e.g.  the  silver-cyanogen  ion  ((Ag(CN)2)~), 
the  silver-ammonia  ion  ((Ag(NH3)2)+),  and  the  copper- 
ammonia  ion  ((Cu(NH3)4)"l"f).  The  ions  formed  by  the  dis- 
sociation of  the  common  compounds  of  the  familiar  ele- 
ments are  shown  in  the  following :  — 


SOLUTION 


147 


TABLE  OP  COMMON  IONS 


El.KMKNT 
OR  GROUP 

ION 

ELKMKNT 
OR  GROUP 

ION 

Kl.KMKNT 

OR  GROUP 

ION 

Hydrogen  .  . 

11+ 

Calcium    .  . 

Ca++ 

Aluminium  . 

A1+++ 

Sodium    .  .  . 

Na+ 

Barium  .  .  . 

Ba++ 

Antimony    . 

SD+++ 

Potassium  .  . 

K+ 

Copper  .  .  . 

Cu++ 

Bismuth    .  . 

B1+++ 

Silver   .... 

Ag+ 

Zinc    .... 

Zn++ 

Iron  (ic)  .  . 

Fe+++ 

Ammonium  . 

NII4+ 

Magnesium  . 

Mg++ 

Tin  (ic)  .  .  . 

Sn++++ 

Mercury  (ous) 

IIg+ 

Lead  .... 

Pb++ 

Chlorine  .   .  . 

01- 

Iron  (ous)    . 

Fe++ 

Bromine  .  .  . 

Br- 

Mercury  (ic) 

Hg++ 

Iodine  .... 

I- 

Tin  (ous)    . 

Sn++ 

Nitrate.  .  .  . 

NO3- 

Sulphate  .  . 

S04— 

Chlorate  .  .  . 

C103- 

Sulphide  .  . 

s— 

Hydroxyl  .  . 

OH- 

Carbonate  . 

co»— 

Chromate    . 

Cr(V 

Dichromate 

Cr,07— 

Many  deductions  arising  from  this  rather  compact  table  will 
be  considered  in  the  succeeding  pages. 

Summary.  —  The  salient  points  discussed  in  this  chapter 
may  be  summarized  as  follows :  The  properties  of  a  solution 
are  mainly  dependent  upon  the  solute  and  its  condition  in 
the  solvent.  Such  properties  as  osmotic  pressure,  freezing 
point,  and  boiling  point,  are  influenced  by  the  number  of 
independent  particles  present  in  the  solution.  When  these 
three  properties  are  measured  independently  in  a  given  solu- 
tion, the  values  agree.  But  when  the  values  are  compared 
under  parallel  conditions  of  measurement,  it  is  found  that 
electrolytes  in  solution  yield  a  larger  number  of  independent, 
individual  particles  than  non-electrolytes;  i.e.  in  electro- 
lytic solutions  some  of  the  molecules  of  the  electrolytes  disso- 
ciate into  ions.  These  ions  are  electrically  charged  atoms 
or  atomic  groups.  Physically  they  act  much  like  molecules. 


148  INORGANIC  CHEMISTRY 

Furthermore,  solutions  of  electrolytes  differ  from  solutions 
of  non-electrolytes  in  conducting  electricity  and  exhibiting 
marked  chemical  activity.  A  study  of  these  two  character- 
istics confirms  the  assumption  of  the  existence  of  electrically 
charged  and  chemically  active  particles  in  the  solution. 

Conclusion.  — The  theory  of  electrolytic  dissociation  as 
outlined  in  the  foregoing  pages  is  not  an  adequate  explana- 
tion of  all  the  facts  of  solution.  It  applies  chiefly  to  dilute 
aqueous  solutions  of  three  classes  of  substances.  Doubtless 
the  present  form  of  the  theory  will  sometime  be  modified  to 
cover  certain  facts  not  at  present  within  its  scope. 

EXERCISES 

1.  Write  out  the  formulas  of  the  ions  formed  when  the  following 
compounds  are  dissolved  separately  in  considerable  water :  Potassium 
chloride,  silver  nitrate,  sodium  chlorate,  ammonium  sulphate,  copper 
nitrate,    calcium    chloride,    zinc    sulphate,    potassium    dichromate, 
calcium  hydroxide. 

2.  Write  the  equations  for  the  following  by  applying  the  method 
for  making  equations  outlined  in  Chapter  VII :    (a)  Iron  and  sulphur 
combine  in  the  ratio  of  7  to  4.     (6)  Ammonia  gas  and  hydrochloric 
acid  gas  form  ammonium  chloride,     (c)  Magnesium  and  hydrochloric 
acid  form  hydrogen  and  magnesium  chloride. 

3.  Discuss :    (a)  Electrolytes  depress  the  freezing  point  abnor- 
mally ;    (6)  ions  migrate  to  their  respective  electrodes. 

4.  Write  the  following  as  ionic  equations :     (a)  Potassium  sul- 
phate and  barium  chloride  form  barium  sulphate  and  potassium 
chloride ;    (6)  sodium  bromide  and  silver  sulphate  form  silver  bro- 
mide and  sodium  sulphate. 


CHAPTER  X 
Acids,  Bases,  and  Salts  —  Neutralization 

Introduction.  —  Many  chemical  compounds  fall  naturally 
into  one  of  three  groups,  long  known  as  acids,  bases,  and  salts. 
Each  group  has  its  characteristic  properties,  though  the 
groups  are  closely  related  and  sometimes  overlap.  Many 
familiar  substances  belong  to  these  groups. 

Acids. — The  common  acids  are  sulphuric  acid  (H2SO4), 
hydrochloric  acid  (HC1),  nitric  acid  (HNO3),  and  acetic  acid 
(C2H4O2).  Many  acids  are  liquid,  as  sulphuric  and  nitric; 
a  few  are  gases,  as  hydrochloric;  others  are  solid,  as  tartaric, 
citric,  oxalic.  Most  acids  are  rather  soluble  in  water,  and 
such  solutions  are  popularly  called  acids.  These  solutions 
may  be  dilute  or  concentrated,  and  the  general  properties 
vary  somewhat  with  the  concentration.  Concentrated  acids 
are  usually  corrosive  and  should  be  handled  with  caution, 
even  when  one  is  familiar  with  their  properties. 

Many  familiar  substances  are  acids  or  contain  them. 
Vinegar,  pickles,  and  similar  relishes  contain  dilute  acetic 
acid.  Lemon  juice  is  mainly  citric  acid.  Sour  milk  con- 
tains lactic  acid.  Unripe  fruits,  sour  bread,  and  sour  wines 
contain  acids.  "Soda  water"  is  a  solution  of  carbonic  acid 
(or  more  accurately  carbon  dioxide),  and  "acid  phosphate" 
is  a  solution  of  a  sour  calcium  phosphate. 

Properties  of  Acids.  —  (1)  Acids,  if  dissolved  in  water, 
usually  have  a  sour  taste.  The  early  chemists  detected  this 
fact,  and  the  term  acid  (from  the  Latin  word  acidus, 

149 


150  INORGANIC  CHEMISTRY 

sour)  emphasizes  this  property.  (2)  Solutions  of  acids 
redden  the  coloring  matter  called  litmus.  Solutions  which 
act  thus  on  blue  litmus  are  described  as  acid,  as  containing 
an  acid,  or  as  having  an  acid  reaction.  (3)  Most  acids 
liberate  free  hydrogen  gas  when  their  solutions  interact  with 
metals.  (4)  Solutions  of  acids  conduct  electricity. 

Composition  of  Acids.  —  All  acids  contain  hydrogen, 
which  is  liberated  in  the  free  state  when  certain  metals  and 
acids  interact.  Most  acids  contain  oxygen.  For  many 
years  it  was  thought  that  oxygen  was  an  essential  component 
of  all  acids,  and  the  name  oxygen  (derived  from  Greek 
words  meaning  "acid  producer")  was  given  to  this  element 
by  Lavoisier  because  of  this  belief.  (See  Discovery  of  Oxygen.) 
We  know  now  that  hydrogen,  not  oxygen,  is  the  essential 
constituent  of  all  acids.  Another  necessary  constituent  of 
acids  is  a  non-metallic  element  like  nitrogen  or  sulphur. 
For  this  reason  it  is  sometimes  convenient  to  think  of  non- 
metals  as  the  elements  which  form  acids.  Thus,  sulphuric 
acid  contains  sulphur,  besides  hydrogen  and  oxygen;  while 
hydrochloric  acid  contains  only  chlorine,  besides  hydrogen. 
The  important  non-metals  which  form  familiar  acids  are 
boron,  carbon,  silicon,  nitrogen,  phosphorus,  sulphur,  flu- 
orine, chlorine,  bromine,  and  iodine. 

Definition  of  an  Acid.  —  For  many  years  an  acid  was 
defined  as  a  compound  producing  a  sour  solution  which 
reddens  blue  litmus,  or  as  a  compound  which  interacts 
chemically  with  a  base,  thereby  forming  a  salt,  or  as  a 
compound  containing  hydrogen  which  can  be  replaced  by  a 
metal.  These  definitions  emphasize  certain  properties  of 
acids,  but  they  are  not  inclusive.  According  to  the  theory 
of  electrolytic  dissociation,  an  acid  is  a  compound  whose 
solution  contains  hydrogen  ions  (H+).  The  sour  taste, 
behavior  toward  litmus,  and  liberation  of  hydrogen  are  due 


ACIDS,   BASES,   AND   SALTS  151 

to  the  hydrogen  ions  which  are  common  to  all  solutions  of 

acids. 

Bases.  —The  common  bases  are  sodium  hydroxide 
(NaOH),  potassium  hydroxide  (KOH),  ammonium  hydrox- 
ide (NH4OH),  and  calcium  hydroxide  (Ca(OH)2).  They  are 
soluble  in  water,  and  such  solutions  are  called  bases;  solu- 
tions of  the  very  soluble  bases  (sodium  hydroxide  and  po- 
tassium hydroxide)  are  often  called  alkalies.  Alkalies,  like 
concentrated  acids,  are  corrosive,  and  should  be  handled 
carefully.  Concentrated  solutions  of  sodium  and  potas- 
sium hydroxides  are  very  corrosive,  and  for  this  reason  are 
called  caustic  alkalies  (from  the  Latin  causticus,  burning). 

Bases  are  components  of  familiar  substances.  Thus, 
ammonia  is  a  solution  of  ammonium  hydroxide.  Lime- 
water  and  baryta  water  are  solutions  of  the  sparingly  soluble 
bases  calcium  hydroxide  and  barium  hydroxide  (Ba(OH)2) 
respectively.  Lye  is  a  concentrated  solution  of  sodium 
hydroxide  or  potassium  hydroxide  (or  both). 

Properties  of  Bases.  —  (1)  Strong,  soluble  bases  have  a 
bitter,  often  biting,  taste;  many,  especially  the  very  soluble 
ones,  have  a  slippery  feeling.  (2)  Soluble  bases  turn  red 
litmus  blue.  Substances  which  act  thus  on  red  litmus  are 
described  as  basic  or  alkaline,  as  having  an  alkaline  reaction, 
or  as  containing  a  base.  (3)  Solutions  of  bases  conduct 
electricity. 

Composition  of  Bases.  —  All  bases  contain  hydrogen  and 
oxygen.  They  also  contain  a  metal,  such  as  sodium.  The 
hydrogen  and  oxygen  are  the  invariable  constituents.  But 
it  is  often  convenient  to  regard  metals  as  the  elements  which 
form  bases,  just  as  the  non-metals  form  acids.  Thus,  the 
base  sodium  hydroxide  is  a  compound  of  the  metal  sodium 
with  hydrogen  and  oxygen.  The  important  metals  which 


152  INORGANIC   CHEMISTRY 

form  familiar  bases  are  sodium,  potassium,  calcium,  and 
barium. 

Definition  of  a  Base.  —  A  base  was  formerly  defined  as 
any  compound  which  has  a  bitter  taste,  turns  red  litmus  blue, 
and  interacts  chemically  with  an  acid,  thereby  forming  a 
salt.  This  definition  emphasizes  certain  properties  of  a 
base,  but  it  is  defective.  According  to  the  theory  of  elec- 
trolytic dissociation,  a  base  is  a  compound  whose  solution 
contains  hydroxyl  ions  (OH~).  The  metal  is  not  the  con- 
stituent which  gives  a  base  its  characteristic  properties. 
These  are  due  to  the  hydroxyl  ions  which  are  common  to  all 
solutions  of  bases. 

Salts.  —  This  is  a  large  and  varied  class  of  compounds. 
The  most  familiar  member  is  sodium  chloride  (NaCl).  It 
is  common  salt  or  table  salt,  and  has  been  known  for  ages. 
Doubtless  this  class  of  chemical  compounds  received  its 
name  from  the  general  resemblance  many  of  them  bear  to 
common  salt.  Most  salts  are  solids  and  are  soluble  in 
water,  although  the  solubility  varies  between  wide  limits. 

Properties  of  Salts.  —  (1)  Salts  often  have  the  well-known 
salty  taste,  though  some  are  bitter,  others  are  astringent,  and 
a  few  have  no  characteristic  taste.  Certain  salts  have  a 
sour  taste.  (2)  Salts  do  not  act  uniformly  on  litmus. 
Some  turn  the  red  to  blue,  others  turn  the  blue  to  red,  and 
many  have  no  action  whatever  on  litmus.  Those  salts 
whose  solutions  do  not  respond  to  the  litmus  test  are  said  to 
be  neutral  or  to  have  a  neutral  reaction.  This  indifference 
to  litmus  is  not  a  decisive  test  for  a  salt,  since  many  other 
substances,  water  for  example,  have  no  action  on  litmus. 
The  term  neutral  is  applied  to  substances  which  do  not 
change  the  color  of  litmus,  whether  or  not  they  are  salts  or 
contain  salts.  (3)  Solutions  of  salts  conduct  electricity. 


ACIDS,   BASES,   AND   SALTS  153 

Composition  of  Salts.  —  Salts  contain  invariably  a  metal 
and  a  non-metal  (which  is  not  hydrogen  or  oxygen).  Most 
salts  also  contain  oxygen.  Thus,  potassium  nitrate  contains 
the  metal  potassium  and  the  non-metal  nitrogen,  besides 
oxygen;  while  potassium  chloride  contains  potassium  and 
the  non-metal  chlorine,  but  no  oxygen.  A  few  salts  contain 
hydrogen  besides  the  characteristic  metal  and  non-metal. 
Thus,  sodium  bicarbonate  contains  hydrogen  besides  the 
metal  sodium,  the  non-metal  carbon,  and  the  non-metal 
oxygen.  These  rather  general  statements  indicate  the  great 
variety  of  salts.  Salts  will  soon  be  further  discussed. 

Neutralization.  — The  nature  and  interrelation  of  acids, 
bases,  and  salts  are  shown  clearly  by  their  chemical  relations. 
When  a  solution  of  an  acid  and  a  base  are  mixed  in  the 
proper  proportion,  they  interact  completely.  The  final 
solution  has  none  of  the  properties  of  an  acid  or  a  base,  but 
it  has  the  properties  characteristic  of  a  salt.  That  is,  the 
acid  and  base  destroy  more  or  less  completely  the  marked 
properties  of  each  other  and  produce  a  salt,  the  latter  being  a 
compound  which  has  few,  if  any,  of  the  properties  of  the 
original  acid  and  base.  The  acid  and  base  neutralize  each 
other.  An  illustration  will  make  this  point  clear.  When 
hydrochloric  acid  and  sodium  hydroxide  interact,  sodium 
chloride  and  water  are  formed.  The  chemical  change  can 
be  written  thus  :  — 

HC1         -f         NaOH       =        NaCl       +       H2O 

Hydrochloric  Sodium  Sodium  Water 

Acid  Hydroxide  Chloride 

This  equation  represents  the  facts  which  have  been  repeatedly 
verified  by  experiment.  The  chemical  change  in  which  an 
acid  and  a  base  neutralize  each  other  and  form  a  salt  and 
water  is  called  neutralization.  Taking  this  equation  as  a 
type  of  the  chemical  change  which  occurs  in  neutralization, 


154  INORGANIC  CHEMISTRY 

it  is  clear  that  (1)  the  metal  of  the  base  takes  the  place  of 
the  hydrogen  of  the  acid,  thereby  forming  a  salt,  while 
(2)  the  hydrogen  of  the  acid  combines  with  the  hydrogen 
and  oxygen  of  the  base  to  form  water.  In  neutralization 
the  hydrogen  and  oxygen  of  the  base  act  as  a  unit.  This 
group  of  atoms  (OH),  as  already  stated,  is  called  hydroxyl. 
Hydroxyl  does  not  exist  free  and  uncombined  like  elements 
and  compounds,  but  it  acts  like  a  single  atom  in  many 
chemical  changes.  It  is  called  a  radical. 

Neutralization  illustrates  double  decomposition.  In  the 
chemical  change  just  cited  both  the  hydrochloric  acid  and 
the  sodium  hydroxide  are  decomposed  and  their  parts  are 
recombined  in  a  different  way;  i.e.  sodium  chloride  and  water 
are  the  new  compounds  resulting  from  the  recombination. 

Neutralization  when  interpreted  by  the  theory  of  elec- 
trolytic dissociation  is  really  the  union  of  hydrogen  ions  with 
hydroxyl  ions.  Suppose  solutions  of  hydrochloric  acid  and 
sodium  hydroxide  are  mixed  in  the  proper  proportions.  The 
mixture  at  first  contains  ions  of  hydrogen,  chlorine,  sodium) 
and  hydroxyl.  But  the  ions  of  hydrogen  and  of  Tiydroxyl 
immediately  unite  to  form  molecules  of  water,  because  water 
does  not  dissociate  into  ions  to  any  appreciable  extent.  The 
final  solution  is  neutral,  because  it  contains  only  ions  of 
sodium  and  chlorine,  the  acid  ions  (H+)  and  the  basic  ions 
(OH~)  having  been  removed  by  their  combination  into 
molecules  of  water.  The  equation  for  the  mutual  neutraliza- 
tion of  hydrochloric  acid  and  sodium  hydroxide  might  be 
written  as  an  ionic  equation,  thus :  — - 

H+  +  Cl-  +  Na+  +  OH-  ->C1-  +  Na+  +  H2O 

The  ionic  equations  for  the  mutual  neutralization  of  other 
pairs  of  acids  and  bases  are  similarly  written.  In  the  case 
just  described,  the  ions  of  sodium  and  of  chlorine  remain 
uncombined  until  the  solution  is  evaporated;  but  as  the  con- 


ACIDS,   BASES,   AND   SALTS 


155 


centration  increases,  the  ions  unite  and  form  molecules  of 
sodium  chloride.  The  latter,  as  already  stated,  is  one  type 
of  the  varied  class  of  compounds  called  salts.  The  salts 
resulting  from  the  combination  of  ions  can  be  obtained  as 
solids  by  the  usual  processes  of  evaporation  to  dryness  or  by 
crystallization.  It  is  evident,  therefore,  that  neutralization 
in  a  broad  sense  is  the  mutual  destruction  of  an  acid  and  a 
base  which  results  in  the  formation  of  a  salt  and  water.  In 
a  narrow  sense  it  is  the  formation  of  molecules  of  water  from 
the  hydrogen  ions  of  the  acid  and  the  hydroxyl  ions  of  the 
base.  This  latter  interpretation  is  supported  by  experi- 
mental evidence.  Heat  is  liberated  when  these  ions  unite  to 
form  water.  If  neutralization  is  merely  the  combination  of 
ions  of  hydrogen  with  ions  of  hydroxyl, 
then  the  same  amount  of  heat  should  be 
liberated  when  a  given  weight  of  water 
is  formed,  whether  the  ions  come  from 
hydrochloric  acid  and  sodium  hydroxide 
or  from  any  other 
pair  of  acid  and 
base.  Experi- 
ment shows  that 
the  heat  of  neu- 
tralization, as  it 
is  called,  is  the 
same  in  all  cases 
of  neutralization, 
when  the  solu- 
tions are  dilute 
and  other  thermal 
changes  do  not 
occur.  It  is  expressed  in  terms  of  a  unit  called  the  calorie,  and 
when  18  gm.  of  water  are  formed  by  the  act  of  neutralization, 
13,700  calories  are  liberated.  (See  Calorie,  Chapter  XI.) 


11- 


12- 


~~ III 


FIG.  18.  —  Burettes.  Enlarged  section  (on  the  left) 
shows  graduations  and  curved  surface  of  the  solu- 
tion, called  the  meniscus.  Correct  reading  is  along 
line  I. 


156  INORGANIC  CHEMISTRY 

Neutralization  is  frequently  brought  about  by  using  bu- 
rettes (Fig.  18).  These  are  graduated  tubes  provided  with 
a  stopcock  to  regulate  the  exit  of  the  solution.  When  prop- 
erly filled,  one  with  an  acid  (e.g.  sulphuric)  and  the  other  with 
a  base  (e.g.  sodium  hydroxide),  a  measured  volume  of  the 
acid  is  allowed  to  flow  into  a  beaker  and  a  few  drops  of  lit- 
mus solution  are  added.  The  solution  of  course  turns  red. 
The  sodium  hydroxide  solution  is  allowed  to  drop  in  slowly 
and  the  mixture  is  stirred  with  a  glass  rod.  As  long  as  an 
excess  of  acid  is  present,  the  color  remains  red.  After 
a  time,  however,  the  color  becomes  purple,  and  an  ad- 
ditional drop  of  sodium  hydroxide  turns  the  solution  blue, 
showing  that  the  acid  has  been  neutralized  by  the  base. 
The  volume  of  sodium  hydroxide  used  is  noted.  If  solutions 
of  known  strength  are  used,  then  the  weights  (found  from  the 
concentration  of  the  solutions)  used  will  be  in  the  same  ratio 
as  the  weights  in  the  equation:  — 

H2SO4  +  2  NaOH  =  Na2SO4  +  2  H2O 

Sulphuric          Sodium  Sodium  Water 

Acid  Hydroxide       Sulphate  36 

98  80  142 

If  the  concentration  of  the  sodium  hydroxide  solution  is 
unknown,  it  can  be  found  by  the  proper  proportion,  because 
the  weights  involved  in  the  chemical  change  are  always  in  the 
ratio  given  in  the  corresponding  equation.  Neutralization 
when  conducted  by  means  of  accurate  apparatus  and  certain 
solutions  of  known  strength  is  an  efficient  method  of  quan- 
titative analysis  and  is  one  of  a  class  called  volumetric 
methods. 

Classification  of  Salts.  —  It  will  be  recalled  that  salts  have 
no  distinctive  class  property  like  acids  and  bases,  such  as 
the  taste  and  behavior  with  litmus.  From  the  standpoint 
of  the  chemical  change  which  occurs  in  complete  neutral- 


ACIDS,  BASES,  AND  SALTS  157 

ization  and  in  analogous  cases  where  the  chemical  change  is 
not  complete,  a  salt  is  a  compound  formed  (1)  by  the  substi- 
tution of  a  metal  for  all  or  part  of  the  hydrogen  of  an  acid, 
or  (2)  by  the  substitution  of  a  non-metal  (or  non-metallic 
group  like  SO4  or  NO3)  for  all  or  some  of  the  hydroxyl  groups 
of  a  base.  There  are,  therefore,  three  classes  of  salts,  — 
normal,  acid,  and  basic.  They  are  prepared  in  various  ways, 
but  it  is  convenient  to  regard  them  as  having,  been  produced 
from  an  acid  or  a  base  by  substitution.  Salts  formed  by  re- 
placing all  the  hydrogen  of  an  acid  by  a  metal  are  called 
normal  salts,  e.g.  sodium  sulphate,  Na2SO4.  On  the  other 
hand,  salts  formed  by  replacing  only  part  of  the  hydrogen 
of  an  acid  by  a  metal  are  called  acid  salts.  Thus,  acid  sodium 
sulphate  (HNaSO4)  may  be  regarded  as  derived  from  sul- 
phuric acid  by  replacing  only  one  of  the  atoms  of  hydrogen 
by  one  atom  of  sodium,  though  of  course  the  salt  is  not  pre- 
pared in  this  way. 

Only  the  acids  containing  two  or  more  replaceable  atoms 
of  hydrogen  can,  as  a  rule,  form  acid  salts ;  e.g.  sulphuric 
acid  (H2SO4)  and  phosphoric  acid  (H3PO4).  Acids  are  often 
classified  by  the  number  of  their  hydrogen  atoms  which  can 
be  replaced  by  a  metal.  This  varying  power  of  replaceability 
is  called  basicity.  A  monobasic  acid  contains  only  one  atom 
of  replaceable  hydrogen  in  a  molecule ;  e.g.  nitric  acid,  HNO3. 
A  molecule  of  acetic  acid  (C2H4O2)  contains  four  atoms  of  hy- 
drogen, but  for  reasons  which  are  too  complex  to  state  here, 
only  one  of  these  atoms  can  be  replaced  by  a  metal ;  it  is  there- 
fore monobasic.  Dibasic  and  tribasic  acids  are  those  that  con- 
tain respectively  two  and  three  replaceable  hydrogen  atoms ; 
e.g.  sulphuric  acid  (H2SO4)  and  phosphoric  acid  (H3P04). 
Normal  salts  may  also  be  regarded  as  formed  by  the  replace- 
ment of  all  the  hydroxyl  groups  of  a  base  by  non-metallic 
atoms  or  atomic  groups.  Thus,  bismuth  nitrate  (Bi(NO3)3) 
is  a  normal  salt  and  may  be  regarded  as  formed  by  the  sub- 


158 


INORGANIC  CHEMISTRY 


stitution  of  three  N03-groups  for  the  three  hydroxyl  groups 
in  bismuth  hydroxide  (Bi(OH)3),  while  basic  bismuth  nitrate 
(Bi(OH)2NO3)  —  a  basic  salt  —  may  be  regarded  as  formed 
by  the  substitution  of  one  N03-group  for  one  of  the  three 
OH-groups  in  the  hydroxide. 

Bases,  like  acids,  are  classified  according  to  their  varying 
power  of  replaceability.  This  power  is  called  acidity.  Bases 
are  called  monacid,  diacid,  triacid  bases,  etc.,  according  to  the 
number  of  the  replaceable  hydroxyl  groups  present  in  a  mole- 
cule. Thus,  sodium  hydroxide  (NaOH)  is  a  monacid  base, 
calcium  hydroxide  (Ca(OH)2)  is  a  diacid  base,  and  aluminium 
hydroxide  (A1(OH)3)  is  a  triacid  base.  Only  bases  having 
two  or  more  replaceable  hydroxyl  groups  form  basic  salts. 

The  relations  of  acids,  bases,  and  salts  may  be  represented 
by  the  following  scheme  :  — 


Acid 
H2S04 

Sulphuric  Acid 


Base 

Zn(OH)2 

Zinc  Hydroxide 


^Normal  S 
Na2S(\ 

Sodium  Sulphate 

ZnCl2 

Zinc  Chloride 


Acid  Salt 
HNaSO4 

Acid  Sodium  Sulphate 


Basic  Salt 
Zn(OH)Cl 

Basic  Zinc  Chloride 


Preparation  of  Salts.  —  It  must  not  be  concluded  from 
the  foregoing  discussion  of  the  kinds  of  salts  that  they  are 
always  prepared  in  the  laboratory  by  mixing  acids  with 
bases  or  metals.  It  is  only  necessary  to  provide  the  metallic 


ACIDS,   BASES,   AND   SALTS  159 

or  non-metallic  constituent  chemically,  so  to  speak.  Salts 
can  be  prepared  in  several  ways.  The  interaction  of  an  acid 
and  a  base  has  been  mentioned.  The  interaction  of  acids 
with  oxides  of  certain  metals  produces  salts.  Sodium  oxide 
and  sulphuric  acid  interact  and  form  the  salt  sodium  sulphate, 
thus:  — 

Na2O     +     H2S04     =     Na2SO4     +     H2O 

Sodium  Sulphuric  Sodium  Water 

Oxide  Acid  Sulphate 

A  metal  and  an  acid  act  similarly.  Zinc  and  sulphuric  acid, 
as  already  stated,  form  the  salt  zinc  sulphate,  thus:  — 


Zn  +     H2S04      =     ZnS04     +  '    2H 

Zinc          Sulphuric  Zinc  Hydrogen 

Acid  Sulphate 

Carbonates  interact  with  acids  and  form  other  salts.  Cal- 
cium carbonate  and  hydrochloric  acid  form  the  salt  calcium 
chloride,  thus:  — 

CaCO3   +      2  HC1     =  CaCl2  +    CO2    +  H2O 

Calcium         Hydrochloric       Calcium        Carbon        Water 
Carbonate  Acid  Chloride       Dioxide 

Sometimes  two  salts  interact  in  solution  and  form  other 
salts  by  double  decomposition.  Sodium  chloride  and  silver 
nitrate  form  the  salts  silver  chloride  and  sodium  nitrate, 
thus:  — 

NaCl        +     AgNO3    =       AgCl  .     +     NaNO3 

Sodium  Silver  Silver  Sodium 

Chloride  Nitrate  Chloride  Nitrate 

Salts  interact  with  certain  acids  and  form  salts  and  other 
acids.  Sodium  chloride  and  sulphuric  acid  form  hydro- 
chloric acid  and  the  salt  sodium  sulphate,  thus:  — 

2  NaCl  +  HsSO*  =      2  HC1     +  Na£O4 

Sodium         Sulphuric       Hydrochloric          Sodium 
Chloride  Acid  Acid  Sulphate 


160 


INORGANIC  CHEMISTRY 


Nomenclature  of  Acids.  —  Oxygen  is  a  component  of  most 
acids,  and  the  names  of  these  acids  correspond  to  the  pro- 
portion of  oxygen  which  they  contain.  The  best  known 
acid  of  an  element  usually  has  the  suffix  -ic;  e.g.  sulphuric 
(HaSO,),  nitric  (HNO3),  phosphoric  (H3PO4).  If  an  element 
forms  another  acid,  containing  less  oxygen,  this  acid  has  the 
suffix  -ous;  e.g.  sulphurous  (H2SOS),  nitrous  (HNO2),  phos- 
phorous (H8PO3).  Some  elements  form  an  acid  containing 
less  oxygen  than  the  -ous  acid;  these  acids  retain  the  suffix 
-ous,  and  have  in  addition  the  prefix  hypo-;  e.g.  hyposulphu- 
rous  (HSO2) ,  hypophosphorous  (H3PO2) ,  hypochlorous  (HC1O) . 
Hypo-  means  under  or  lesser.  If  an  element  forms  an  acid 
containing  more  oxygen  than  the  -ic  acid,  such  an  acid  re- 
tains the  suffix  -ic,  and  has  in  addition  the  prefix  per-;  e.g. 
persulphuric  (H2S208),  perchloric  (HC1O4).  The  prefix  per- 
means  beyond  or  over.  The  few  acids  which  contain  no 
oxygen  have  the  prefix  hydro-  and  the  suffix  -ic;  e.g.  hydro- 
chloric (HC1),  hydrobromic  (HBr),  hydrofluoric  (HF),  and 
hydriodic  (HI).  It  should  be  noticed  that  the  suffixes  ic- 
and  -ous  are  not  always  added  to  the  whole  name  of  the 
element,  but  often  to  some  modification  of  it. 

The  nomenclature  of  acids  is  well  illustrated  by  the  series 
of  chlorine  acids:  — 

ACIDS  OF  THE  ELEMENT  CHLORINE 


NAMK 

FORMULA 

HC1 

HC1O 

HC1O2 

Chloric  .     .          .          

HClOg 

Perchloric        .                              ... 

HC1O4 

Not  all  elements  form  a  complete  series  of  acids,  but  the 
nomenclature  usually  agrees  with  the  above  principles. 


ACIDS,   BASES,   AND   SALTS  161 

Some  acids  have  commercial  names.  Thus,  sulphuric 
acid  is  often  called  oil  of  vitriol,  and  hydrochloric  acid  is 
known  as  muriatic  acid.  Acids  in  which  carbon  is  an  essen- 
tial constituent  end  in  -ic,  but  they  are  often  arbitrarily 
named.  (See  Organic  Acids.) 

Nomenclature  of  Bases.  —  There  is  no  general  rule  for  the 
nomenclature  of  bases,  as  in  the  case  of  acids.  Since  most 
bases  contain  hydrogen  and  oxygen,  they  are  called  hydrox- 
ides. The  term  hydrate  is  sometimes  used  as  a  synonym 
of  hydroxide,  but  it  more  correctly  describes  those  com- 
pounds in  which  water  is  one  constituent.  Thus,  crystal- 
lized salts  containing  water  of  crystallization  are  often 
called  hydrates.  The  term  alkali  emphasizes  general  prop- 
erties rather  than  suggests  specific  composition  and  is  now 
applied  to  the  very  active  bases  such  as  sodium  and  potas- 
sium hydroxides.  Hydroxides  are  distinguished  from  each 
other  by  placing  the  name  of  the  metal  before  the  word 
hydroxide;  e.g.  sodium  hydroxide,  potassium  hydroxide, 
calcium  hydroxide. 

The  common  hydroxides  have  long  been  known  by  special 
names.  Thus,  a  solution  of  calcium  hydroxide  is  sometimes 
called  limewater.  Ammonium  hydroxide  solution  is  some- 
times called  ammonia  water  or  simply  ammonia,  and  it  was 
formerly  called  volatile  alkali.  The  hydroxides  of  sodium 
and  potassium  are  often  called  caustic  soda  and  caustic 
potash,  and  occasionally  the  term  fixed  alkali  is  used  to 
emphasize  the  fact  that  they  are  non-volatile. 

Nomenclature  of  Salts.  —  The  names  of  salts  containing 
oxygen  are  derived  from  the  name  of  the  corresponding  acid. 
The  characteristic  suffix  of  the  acid  is  changed  to  indicate 
this  relation.  Thus,  the  suffix  -ic  becomes  -ate,  and  the  suffix 
-ous  becomes  -ite.  Hence  :  -=- 


162  INORGANIC  CHEMISTRY 

Sulphuric  acid  forms  sulphates. 
Sulphurous  acid  forms  sulphites. 
Nitric  acid  forms  nitrates. 
Nitrous  acid  forms  nitrites. 
Chloric  acid  forms  chlorates. 
Hypochlorous  acid  forms  hypochlorites. 
Permanganic  acid  forms  permanganates. 

The  name  of  the  replacing  metal  is  retained ;  e.g.  potassium 
chlorate,  sodium  sulphate,  calcium  hypochlorite,  potassium 
permanganate.  Notice  that  the  prefixes  hypo-  and  per-  are 
not  changed. 

The  names  of  salts  containing  only  two  elements,  follow- 
ing the  general  rule  for  binary  compounds,  end  in  -ide.  This 
suffix  is  added  to  a  modification  of  the  name  of  the  non-metal, 
giving  the  names  chloride,  bromide,  sulphide,  fluoride,  etc. 
The  prefix  hydro-  which  is  contained  in  the  name  of  the  acid 
is  omitted.  Thus,  the  sodium  salt  of  hydrochloric  acid  is 
sodium  chloride ;  similarly,  there  are  the  names  potassium 
chloride,  calcium  fluoride,  and  sodium  iodide. 

Relation  of  Oxides  to  Acids  and  Bases.  —  Most  non- 
metallic  elements  form  oxides  which  unite  with  water  and 
produce  an  acid.  The  oxides  of  many  metallic  elements,  on 
the  other  hand,  unite  with  water  and  produce  bases.  The 
two  oxides  of  the  non-metal  sulphur  act  thus :  — 

502  +      H2O     =     H2S03 

Sulphur  Water  Sulphurous 

Dioxide  Acid 

503  +      H20     =     H2SO4 

Sulphur  Water  Sulphuric 

Trioxide  Acid 

The  oxide  of  the  metal  calcium  acts  thus :  — 
CaO         +      H2O     =     Ca(OH)2 

Calcium  Water  Calcium 

Oxide  Hydroxide 


ACIDS,   BASES,   AND   SALTS  163 

Oxides  of  non-metals  which  unite  with  water  and  produce 
acids  are  called  acid  forming  oxides  or  acid  anhydrides,  i.e. 
literally,  substances  without  water.  Examples  are  carbonic 
anhydride  (C02),  sulphuric  anhydride  (SO3),  phosphoric 
anhydride  (P2O5).  Oxides  of  metals  which  produce  basic 
hydroxides  are  base  forming  oxides  or  basic  anhydrides. 
Examples  are  calcium  oxide  (CaO),  sodium  oxide  (Na2O), 
barium  oxide  (BaO).  A  few  oxides  behave  exceptionally.  It 
is  convenient  to  regard  an  acid  anhydride  as  the  root  or  basis 
of  its  corresponding  acid,  and  a  basic  anhydride  as  the  root 
of  its  hydroxide. 

The  fact  that  many  non-metallic  oxides  redden  moist  blue  litmus 
led  Lavoisier  into  the  erroneous  belief  that  oxygen  is  an  essential 
constituent  of  all  acids.  And  some  authorities  even  now  speak 
(incorrectly)  of  these  oxides  as  acids;  thus,  carbon  dioxide  (COz)  is 
occasionally  called  carbonic  acid.  The  compounds  which  Lavoisier 
called  acids  were  anhydrides.  And  it  was  not  until  about  1811  that 
Davy  showed  (1)  that  some  acids  do  not  contain  oxygen  (e.g.  hydro- 
chloric acid,  HC1),  and  (2)  that  the  so-called  acids  of  Lavoisier  are 
not  real  acids  until  they  obtain  hydrogen  from  the  water  in  which 
they  dissolve. 

Degree  of  Dissociation  of  Acids,  Bases,  and  Salts.  —  The 
degree  to  which  acids,  bases,  and  salts  dissociate  is  due  to 
two  factors,  viz.  the  nature  of  the  substance  itself  and  the 
concentration  of  the  solution.  In  general,  dissociation  is 
slight  in  a  concentrated  solution,  and  increases  as  the  solu- 
tion becomes  more  and  more  dilute.  Thus,  in  concentrated 
nitric  acid  (62  per  cent)  the  per  cent  of  ionized  substance  is 
only  about  9.6,  while  in  a  dilute  solution  (.63  per  cent)  the 
per  cent  is  about  90.  The  different  degrees  of  dissociation 
of  acids,  bases,  and  salts  may  be  readily  seen  by  consulting 
the  table  of  the  dissociation  of  these  substances  on  the  next 
page.  In  order  to  compare  the  varying  degrees  of  dissocia- 
tion, solutions  of  the  same  relative  strength  (at  a  fixed  tem- 
perature) must  be  taken.  Usually  normal  solutions  at  18°  C. 


164 


INORGANIC  CHEMISTRY 


are  selected,  i.e.  solutions  which  contain  in  a  liter  a  number 
of  grams  numerically  equal  to  the  molecular  weight  divided 
by  the  number  of  replaceable  hydrogen  atoms  present  (or 
their  metallic  equivalent).  For  example,  a  normal  solution 
of  hydrochloric  acid  contains  36.5  gm.  to  the  liter,  i.e.  (35.5 
+  l)-s-  1;  but  since  sulphuric  acid  contains  two  replace- 
able hydrogen  atoms,  its  normal  solution  contains  (2  +  32 
H-  64)  -5-  2  =  49  gm.  in  each  liter  of  solution.  Similarly,  a 
normal  solution  of  sodium  hydroxide  contains  (23  +  16  +  1) 
-f-  1  =  40  gm.,  while  one  of  potassium  sulphate  contains 
(78  +  32  +  64)  -T-  2  =  87  gm.  to  the  liter  of  solution.  Nor- 
mal solutions  are  often  designated  by  the  letter  N,  and  differ- 
ent concentrations  are  indicated  correspondingly,  e.g.  tenth 

"NT 

normal  by  — ,   twice  normal  by  2  N.     The   approximate 

per  cent  of  dissociation  of  certain  acids,  bases,  and  salts  is 
tabulated  below;  the  solutions  are  one  tenth  normal  (at  18°C.). 

TABLE  OF  DISSOCIATION  OF  ACIDS,  BASES,  AND  SALTS 


ELECTROLYTE 


PER  CENT  OF  DISSOCIATION 


Nitric  acid  (H+,  N08-) 

Hydrochloric  acid  (H+,  Cl~)  .  .  . 
Sulphuric  acid  (H+,  HSO4-)  .  .  . 
Acetic  acid  (H+,  C2H3O2-)  .  .  . 
Carbonic  acid  (H+,  HCO3-)  .  .  . 
Hydrocyanic  acid  (H+,  CN~).  .  . 
Potassium  hydroxide  (K+,  OH~)  .  . 
Sodium  hydroxide  (Na+,  OH~)  .  . 
Ammonium  hydroxide  (NH4+,  OH~) 
Potassium  chloride  (K+,  Cl~).  .  . 
Sodium  chloride  (Na+,  Cl~)  .  .  . 
Potassium  nitrate  (K+,  NO3-)  .  . 
Potassium  sulphate  (2  K+,  S04-  -) 
Silver  nitrate  ((Ag+,  NO3-)  .  .  . 
Copper  sulphate  (Cu++,  SO4-  -) .  . 


90 
90 
60 

1.3 
.17 
.01 
90 
90 

1.4 
86 
84 
83 
72 
86 
37 


ACIDS,   BASES,   AND   SALTS  165 

An  examination  of  the  table  above  shows  more  or  less 
equality  of  dissociation  among  salts,  but  a  rather  wide 
variation  in  the  case  of.  both  acids  and  bases.  The  term 
strong  is  applied  to  those  acids  and  bases  which  dissociate 
to  a  marked  degree,  and  the  term  weak  to  those  whose  dis- 
sociation is  limited. 

Salts  and  the  Theory  of  Electrolytic  Dissociation.  —  Ac- 
cording to  the  theory  of  electrolytic  dissociation,  solutions  of 
normal  salts  contain  neither  hydrogen  nor  hydroxyl  ions; 
and,  as  a  rule,  they  are  neutral  to  litmus.  But  solutions  of 
acid  salts  contain  the  ions  characteristic  of  an  acid  (i.e. 
H-ions  or  H+)  as  well  as  the  ions  of  a  salt.  While  solutions 
of  basic  salts  contain  the  ions  characteristic  of  a  base  (i.e. 
hydroxyl  ions  or  OH~)  as  well  as  the  ions  of  a  salt.  A  salt 
may  be  denned  in  terms  of  the  theory  of  electrolytic  disso- 
ciation as  the  compound  ultimately  formed  by  the  union  of 
one  or  more  metallic  ions  of  the  base  and  one  or  more  non- 
metallic  ions  of  an  acid,  supplemented  in  some  cases  by  one 
or  more  hydrogen  or  hydroxyl  ions.  In  other  words,  salts 
are  chemical  compounds  which  have  a  certain  composition 
but  not  necessarily  uniform  behavior.  For  example,  take 
the  behavior  toward  litmus.  Normal  salts  are  neutral,  acid, 
or  basic;  acid  salts  are  acid  or  nearly  neutral;  most  basic 
salts  are  nearly  insoluble  in  water,  and  exhibit  a  faint 
reaction  toward  litmus.  The  terms  normal,  acid,  and 
basic  as  applied  to  salts  indicate  their  composition  but 
do  not  describe  their  properties.  An  explanation  of  this 
varying  behavior  of  salts  toward  litmus  is  offered  by  the 
theory  of  electrolytic  dissociation.  Hitherto,  pure  water 
has  been  referred  to  as  a  non-electrolyte,  and  in  most  solu- 
tions there  is  no  evidence  of  its  dissociation  into  the  ions 
H+  and  OH~.  In  some  cases,  however,  the  very  slight 
ionization  of  water  becomes  a  significant  factor  in  establish- 


166  INORGANIC  CHEMISTRY 

ing  the  properties  revealed  by  the  solution;  i.e.  its  ions, 
though  comparatively  very  few  in  number,  interact  with  the 
ions  of  certain  substances  and  thereby  produce  very  interest- 
ing results.  Thus,  sodium  carbonate  yields  the  ions  2  Na+ 
and  C03~ " ;  but  CO3-ions  are  unstable  and  combine  with 
H-ions  to  form  the  ion  HC03~  -  -  the  last  named  being  one 
of  the  ions  yielded  by  the  slightly  dissociated  acid  H2C03. 
The  removal  of  H-ions  leaves  an  excess  of  OH-ions,  which 
give  the  solution  an  alkaline  reaction  toward  litmus.  On 
the  other  hand,  copper  sulphate  (a  normal  salt  from  the 
standpoint  of  composition)  has  an  acid  reaction.  It  yields 
the  ions  Cu++  and  S04~ " ;  but  the  Cu-ions  combine  with 
OH-ions  to  form  copper  hydroxide  (Cu(OH)2),  which  dis- 
sociates only  to  a  very  slight  extent.  The  removal  of  OH-ions 
leaves  an  excess  of  H-ions  in  the  solution  and  gives  it  an 
acid  reaction.  For  similar  reasons  the  litmus  reaction  of 
potassium  carbonate  and  potassium  cyanide  is  alkaline, 
while  the  reaction  of  ferric  chloride  is  acid. 

Acid  and  basic  salts  are  readily  interpreted  by  the  theory. 
Acid  sodium  sulphate  (HNaSO4)  yields  Na+  and  HSO4~, 
but  the  latter  ion  dissociates  to  some  extent  into  H+  and 
S04~  ~ ;  the  solution  is  therefore  rendered  acid  to  litmus  by 
the  H-ions.  But  acid  sodium  carbonate  (HNaC03)  is  nearly 
neutral  to  litmus.  It  yields  the  ions  Na+  and  HCO3~ ;  but 
since  the  latter  dissociates  to  only  a  very  slight  extent,  the 
solution  is  not  provided  with  an  excess  of  either  H-ions  or 
OH-ions,  and  therefore  has  only  a  fairly  neutral  reaction. 
Basic  salts  may  be  similarly  explained. 

Chemical  changes  like  those  cited  in  the  first  paragraph  are 
examples  of  hydrolysis,  i.e.  a  chemical  change  involving 
water  and  certain  salts.  It  is  typically  exhibited  by  salts 
derived  from  strong  bases  (e.g.  NaOH  and  KOH)  and  weak 
acids  (e.g.  H2CO3  and  HCN),  or  from  weak  bases  (e.g. 
Cu(OH)2  and  Fe(OH)3)  and  strong  acids  (e.g.  HC1  and 


ACIDS,   BASES,   AND   SALTS  167 

H2S04).  The  behavior  of  the  final  solution  toward  litmus 
depends  upon  the  composition  of  the  salt,  as  the  following 
equations  show.  The  ordinary  chemical  equation  for  the 
hydrolysis  of  sodium  carbonate  is  - 

Na2CO3   +  H2O   =     HNaCO3  +    NaOH 

Sodium  Water  Acid  Sodium  Sodium 

Carbonate  Carbonate  Hydroxide 

The  ionic  equation  is  — 

2  Na+  +  C03-  "  +  H+  +  OH~  ->2  Na+  +  HCOr  +  OH- 

The  corresponding  equations  for  the  hydrolysis  of  copper 
sulphate  are  — 

CuSO4  +  2  H2O  =  Cu(OH),  +  HaSO, 

Copper  Water  Copper         Sulphuric 

Sulphate  Hydroxide          Acid 

Cu++  +  SOr  ~  +  2  H+  +  2  OH--*Cu(OH),  +  SOr  -  +  2  H+ 


PROBLEMS  AND  EXERCISES 

1.  Name    the    hydroxide    corresponding    to    sodium,    potassium, 
calcium,  barium,  zinc,  lead,  copper. 

2.  Name  the  potassium  salt  of  manganic  acid,   calcium  salt  of 
hydrofluoric  acid,  sodium  salt  of  carbonic  acid,   potassium  salt  of 
tartaric  acid,  lead  salt  of  chromic  acid,  potassium  salt  of  hydrobromic 
acid,  potassium  salt  of  permanganic  acid. 

3.  Name  the  sodium  salt  of 'hydrochloric  acid.     Name  the  cor- 
responding salt  of  potassium,  lead,  calcium,  barium,  zinc,  silver. 

4.  Name  the  same  salts  of  nitric  acid.     Of  nitrous  acid. 

5.  Name  the  same  salts  of  sulphuric  acid.     Of  hypochlorous  acid. 
Of  perchloric  acid. 

6.  Give  the  name  and  formula  of  the  ions,  if  any,  formed  by 
(a)  potassium  sulphate,   (6)  silver  chloride,   (c)  barium  sulphate, 
(d)  lead  nitrate,  (e)  lead  sulphate,  (/)  lead  chromate. 

7.  Classify  into  acids,  bases,  salts,  and  anhydrides:  SO3,  Pb(OH)2, 
HBr,  KC1O4,  CO2,  H8PO4,  P2O6,  NH4C1,  Ag2SO4,  Ca(NO3)2,  HNO2, 
Ba(OH)2,  HI,  HNaC03,  Cu(OH)2,  FeCl3,  HNaSO4,  NaNO2,  SO2,  CaO. 


168  INORGANIC  CHEMISTRY 

8.  What  is  the  percent  of  hydrogen  in  (a)  sulphuric  acid,  (6)  hy- 
drochloric acid,  (c)  nitric  acid  ? 

9.  What  is  the  percent  of  OH  in  (a)  NaOH,  (6)  KOH,  (c)  NH4- 
OH,  (d)  Ca(OH)2? 

10.  Write  equations  for  the  following  reactions  :     (a)  Hydro- 
chloric acid  neutralizes  sodium  hydroxide  and  forms  sodium  chlo- 
ride and  water.     (6)  Nitric  acid  neutralizes  potassium  hydroxide 
and  forms  potassium  nitrate  and  water,      (c)  Potassium  chloride  and 
silver  nitrate  form  silver  chloride  and  potassium  nitrate,      (d)  Car- 
bon dioxide  and  water  form  carbonic  acid  (H2CO3).      (e)  Barium 
oxide  and  water  form  barium  hydroxide.     (/)  Potassium  carbo- 
nate by  hydrolysis  forms  acid  potassium  carbonate  and  potassium 
hydroxide. 

11.  Write  the  ionic  equations  for  (a)  and  (6)  in  Problem  10. 

12.  Express  the  following  reactions  as  ordinary  chemical  equa- 
tions and  as  ionic  equations  :  (a)  Sodium  bromide  and  silver  nitrate 
form  silver  bromide  and  sodium  nitrate;      (6)  ammonium  iodide 
and  silver  sulphate  form  silver  iodide  and  ammonium  sulphate. 

13.  Complete  and  balance  :  (a)  Ag2S04  +  KBr  =  K2SO4  H  ---  ; 
(6)     BaF2  +  H2SO4  =  HF  +  --  ;       (c)     NaBr  :  +  C12  =  Br  +  -  ; 
(d)    H2SH  ---  =  HI  +  -    -;     (e)    HI  +  O  =  H20  +  I  ;     (/)  PBr3 


14.  What  weight  of  sodium  hydroxide  is  needed  to  neutralize 
45  cc.  of  a  sulphuric  acid  solution  having  the  specific  gravity  1.3 
and  containing  32  per  cent  (by  weight)  of  pure  H2SO4? 

15.  In  what  proportions  by  weight  do  the  following  form  a  neu- 
tral solution  :     (a)  Hydrobromic  acid  (HBr)  and  ammonium  hy- 
droxide,   (6)    phosphoric    acid    (H3PO4)    and    calcium    hydroxide, 
(c)  sulphurous  acid  and  sodium  hydroxide? 

16.  Calculate  the  weight  of  the  salt  formed  in  the  following 
cases  of  neutralization  :    (a)  Hydrochloric  acid  and  10  gm.  of  potas- 
sium hydroxide,  (6)  nitric  acid  and  37  gm.  of  barium  hydroxide, 
(c)  sodium  hydroxide  and  55  gm.  of  acetic  acid. 

17.  If  50  gm.  of  sodium  nitrate  (92  per  cent  pure)  is  made  into 
nitric  acid,  what  weight  of  sodium  hydroxide  (92  per  cent  pure) 
will  neutralize  the  acid  ? 

18.  A  ton  of  calcium  acetate  is  needed.     Calculate  the  required 
weight  of  lime  (97  per  cent  CaO)  and  the  volume  of  acetic  acid 
solution  having  the  specific  gravity  1.035  and  containing  25  per 
cent  (by  weight)  of  the  pure  acid. 


CHAPTER  XI 
Energy  and  Chemical  Change  —  Chemical  Equilibrium 

CHEMICAL  change  is  always  attended  by  the  production  or 
consumption  of  one  or  more  of  the  different  forms  of  energy, 
such  as  light,  heat,  and  electricity.  This  means  that  a 
chemical  change  involves  not  only  a  rearrangement  of  matter, 
but  also  a  transformation  or  redistribution  of  energy.  Thus, 
when  coal  is  burned,  a  new  compound,  called  carbon  dioxide, 
is  formed,  but  heat  is  also  liberated  as  a  result  of  the 
chemical  change.  Sometimes  we  pay  more  attention  to 
the  redistributed  matter  than  to  the  energy,  but  both 
are  involved.  In  the  present  chapter  we  shall  emphasize 
the  relation  of  energy  to  chemical  change.  (See  Chapter  I.) 

The  law  of  the  conservation  of  energy  should  be  recalled 
in  this  connection  (see  page  8) .  Energy  cannot  be  created 
or  destroyed  by  any  known  means ;  we  can  only  transform 
it.  Hence,  the  chemical  energy  that  is  in  elements  and 
compounds  appears  as  heat,  light,  or  electricity  when  chemi- 
cal changes  occur.  And  these  forms  of  energy  are  in  turn 
transformed  into  chemical  energy  and  stored  up  in  the 
chemical  elements  and  compounds  which  are  the  result  of 
chemical  changes. 

LIGHT  AND  CHEMICAL  CHANGE 

Light  is  often  produced  by  chemical  change.  Sometimes 
the  light  is  faint,  as  in  the  slow  oxidation  of  yellow  phos- 
phorus, which  is  luminous  in  moist  air.  This  phenomenon 
is  also  exhibited  by  mixtures  containing  even  a  very  little 

169 


170  INORGANIC  CHEMISTRY 

phosphorus ;  for  example,  the  head  of  a  phosphorus  match. 
When  the  phosphorus  and  oxygen  unite,  part  of  the  energy 
in  each  element  is  transformed  into  light  and  part  is  stored 
up  in  the  phosphorus  pentoxide.  Usually  the  transforma- 
tion of  chemical  energy  into  light  is  more  vivid ;  that  is, 
more  chemical  energy  becomes  light.  Many  chemical  ex- 
periments are  accompanied  by  intense  light,  especially  those 
involving  combination  with  oxygen  or  with  chlorine.  Thus, 
magnesium  burns  in  oxygen  with  a  dazzling  light,  and 
powdered  antimony,  as  well  as  some  other  metals,  bursts 
into  a  flame  when  dropped  into  chlorine  gas.  Combustion 
in  general,  especially  of  coal,  oils,  and  gases  containing 
compounds  of  carbon,  is  usually  attended  by  light,  and 
serves  as  an  excellent  illustration  of  the  transformation  of 
chemical  energy  into  light.  (See  Chapter  XVI.) 

Light  is  also  often  transformed  into  chemical  energy. 
This  transformation  is  typically  illustrated  by  photographic 
processes.  Paper,  glass  plates,  and  films  coated  with  com- 
pounds of  silver  are  blackened  on  exposure  to  the  light; 
the  compounds  are  changed  chemically  and  light  is  trans- 
formed into  chemical  energy.  Another  transformation  more 
or  less  familiar  is  that  involved  in  the  fading  of  colored 
fabrics,  wall  paper,  and  paintings.  Light  is  also  absolutely 
essential  in  the  complex  chemical  changes  involved  in  the 
growth  of  plants.  The  sunlight  is  stored  up  in  the  plants 
and  is  subsequently  utilized  by  mankind  when  wood  and 
coal  are  burned  as  fuel  or  vegetable  matter  is  consumed  as 
food.  Certain  chemical  changes  which  proceed  very  slowly 
are  hastened  by  light.  Thus,  hydrogen  and  chlorine  gases 
when  mixed  in  the  dark  do  not  unite  perceptibly,  but  they 
combine  slowly  in  diffused  light  and  instantaneously  in  the 
direct  sunlight.  Similarly,  a  solution  of  chlorine  (in  water) 
evolves  oxygen  slowly  in  the  dark  but  more  rapidly  in  the 
light. 


HEAT  171 

HEAT  AND  CHEMICAL  CHANGE 

Heat  and  chemical  change  are  closely  and  definitely  re- 
lated. A  chemical  change  is  almost  invariably  accompanied 
by  the  liberation  or  absorption  of  heat,  usually  the  libera- 
tion. Vigorous  and  rapid  reactions  develop  considerable 
heat  and  are  also  often  attended  by  light,  while  the  heat 
evolved  by  feeble  or  slow  reactions  is  comparatively  slight 
and  sometimes  can  scarcely  be  detected. 

A  familiar  instance  of  the  evolution  of  heat  by  chemical 
change  is  the  slaking  of  lime.  Lime  is  calcium  oxide  (CaO), 
and  when  lime  and  water  are  mixed  their  chemical  union 
produces  sufficient  heat  to  boil  water  and  often  to  set  fire 
to  wood.  Steam  can  be  seen  escaping  from  the  boxes  in 
which  lime  is  being  mixed  with  water  in  the  preparation  of 
plaster  or  mortar.  Buildings  in  which  lime  is  stored  some- 
times take  fire,  if  rain  leaks  in  upon  the  lime.  Ships  loaded 
with  lime  are  in  constant  danger  of  being  set  afire.  An- 
other illustration  is  provided  by  the  combustion  of  fuels 
such  as  coal,  wood,  oils,  and  gases.  These  are  largely  car- 
bon. The  carbon  in  these  substances  unites  with  the  oxygen 
of  the  air,  and  the  chemical  energy  in  both  elements  becomes 
heat  to  a  great  extent,  some  of  course  remaining  in  the 
products  of  combustion.  Many  chemical  changes,  already 
considered,  are  attended  by  the  liberation  of  heat,  the  most 
conspicuous  being  the  act  of  combination  with  oxygen. 
Thus,  when  hydrogen  burns,  the  act  of  combination  is 
strikingly  manifested  by  the  colorless,  intensely  hot  flame. 
The  interaction  of  metals  and  acids,  as  seen  in  the  prepara- 
tion of  hydrogen,  develops  heat.  The  chemical  union  of 
sulphur  and  metals  is  often  accompanied  by  heat  sufficient 
to  cause  incandescence.  Reduction  of  metallic  oxides  often 
liberates  much  heat.  For  example,  the  heat  attending  the 
reduction  of  iron  oxide  by  aluminium  is  so  intense  that  the 
iron  melts.  (See  Thermit.) 


172  INORGANIC   CHEMISTRY 

Many  chemical  changes  take  place  slowly  at  the  ordinary 
temperature.  Once  started  by  heat,  however,  they  proceed 
until  the  interacting  substances  are  exhausted  or  the  ex- 
ternal supply  of  heat  is  removed.  There  are  many  illus- 
trations of  this  hastening  of  chemical  change  by  heat. 
Magnesium  tarnishes  very  slowly  in  the  air,  but  if  a  lighted 
match  is  applied  to  the  metal,  oxidation  proceeds  rapidly 
until  the  magnesium  is  entirely  changed  into  magnesium 
oxide.  Hydrogen  and  oxygen  mix  freely  at  the  ordinary 
temperature  without  appreciable  combination;  if  mixed  in 
the  proportion  of  two  volumes  to  one  and  heated  to  600°  C., 
combination  takes  place  in  about  one  hour,  while  .union  in- 
stantly occurs  when  heat  is  applied  in  the  form  of  a  flame 
or  an  electric  spark.  Similarly,  illuminating  gas  must  be 
lighted  before  it  will  interact  chemically  with  the  oxygen  of 
the  air,  but  once  raised  to  the  proper  temperature,  the 
chemical  change  continues  as  long  as  the  gas  is  supplied. 
Combustible  substances,  such  as  wood,  many  oils  and  gases, 
sulphur,  and  phosphorus,  must  be  raised  to  a  minimum  tem- 
perature called  the  kindling  temperature  before  the  chemical 
changes  attending  combustion  can  proceed.  The  kindling 
temperature  varies  with  the  physical  condition  in  the  case  of 
many  solids,  being  lower,  as  a  rule,  when  the  substance  is 
finely  divided  or  presents  a  relatively  large  surface.  Thus, 
shavings  catch  fire  at  a  lower  temperature  than  a  stick  or 
log  of  the  same  variety  of  wood.  In  many  chemical  changes 
heat  must  be  constantly  supplied.  Thus,  mercuric  oxide 
decomposes  into  mercury  and  oxygen  as  long  as  it  is  suffi- 
ciently heated,  but  when  the  flame  is  removed,  the  chemical 
change  slackens  and  soon  ceases  altogether.  The  same  is 
true  of  potassium  chlorate.  Indeed,  heat  is  one  of  the 
most  efficient  aids  to  chemical  change,  and  various  sources 
of  heat  are  indispensable  in  the  laboratory  as  well  as  in 
chemical  manufactories. 


HEAT  173 

Sources  of  Heat.  —  Heat  is  so  essential  in  all  chemical 
operations  that  chemists  have  devised  and  used  many  ap- 
pliances for  generating  intense  and  continuous  heat.  The 
alchemists  burned  wood  and  charcoal  in  the  furnaces  which 
heated  their  crucibles.  Priestley  and  Lavoisier  employed  a 
lens  or  burning  glass  in  some  of  their  experiments.  Liebig 
and  his  contemporaries  used  a  charcoal  furnace  in  analyzing 
organic  compounds.  The  greatest  advance  was  made  when 
Bunsen  invented  the  burner  which  consumes  gas  and  pro- 
duces a  hot,  smokeless  flame.  This  burner  is  replaced  by 
the  blast  lamp,  oxyhydrogen  blowpipe,  or  oxyacetylene  blow- 
pipe when  a  high  temperature  is  desired.  In  the  arts  and 
industries  various  kinds  of  furnaces  are  used ;  e.g.  rever- 
beratory,  open  hearth,  and  blast  (see  Figs.  64,  82,  and  80). 
But  all  these  sources  of  heat  have  been  surpassed  by  the 
electric  furnace.  It  is  well 
known  that  an  electric  arc 
light  produces  intense  heat. 
The  high  temperature  of 
the  arc,  i.e.  space  between 
the  glowing  ends  of  the 
carbons,  is  unequaled  by 
that  of  any  other  source 

r         ,.£    .11  T/.    ,,         FIG.  19.  —  Electric  furnace  —  arc  type. 

of  artificial   heat.     If  the 

carbon  rods  are  inclosed  in  an  infusible  box  or  vessel 
that  prevents  escape  of  heat,  a  temperature  estimated  to 
be  about  3500°  C.  can  be  produced  inside  the  receptacle. 
This  apparatus  is  called  an  electric  furnace.  One  type  of 
the  electric  furnace  is  shown  in  Figure  19.  When  a  cur- 
rent is  passed  through  the  carbon  rods,  the  tremendous 
heat  produced  within  the  space  is  retained  by  the  non- 
conducting walls  and  acts  upon  the  substance  in  the  cru- 
cible below  the  arc.  The  outside  of  the  furnace  remains 
cold  enough  to  be  touched  by  the  hand,  but  the  inside 


174  INORGANIC  CHEMISTRY 

is  almost  twice  as  hot  as  the  oxyhydrogen  flame.  There 
is  no  electrical  action  upon  the  chemicals.  The  intense 
heat  alone  produces  the  remarkable  physical  and  chemical 
changes,  and  for  this  reason  the  process  is  often  called  an 
electrothermal  process.  Sand,  lime,  magnesium  oxide,  and 
other  refractory  substances  melt  and  volatilize.  The  ele- 
ments carbon,  silicon,  and  boron  boil;  and  gold,  copper, 
and  platinum  quickly  melt  and  vaporize.  Large  masses  of 
rare  and  uncommon  elements  are  quickly  reduced  from 
their  oxides  and  obtained  in  the  pure  state;  e.g.  chromium, 
manganese,  tungsten,  uranium,  and  molybdenum.  Stable 
compounds  of  carbon,  boron,  and  silicon  are  formed.  These 
are  the  carbides,  borides,  and  silicides.  Some  of  the  car- 
bides have  an  industrial  use  as  well  as  scientific  interest, 
especially  calcium  carbide  and  silicon  carbide. 

Another  type  of  electric  furnace,  known  as  the  resistance 
type,  is  shown  in  Figure  20.  It  is  essentially  an  insulated 
box-like  structure  of 
heat-resisting  mate- 
rials.  The  wires  con- 
veying  the  current  are 
attached  to  perma- 

,  ,      FIG.  20.  —  Electric  furnace  —  resistance  type. 

nent  outer  ends  and 

the  carbon  electrodes  project  into  the  furnace.  Pieces  of 
broken  carbon  make  electrical  connection  between  the  elec- 
trodes and  at  the  same  time  offer  great  resistance  to  the 
current.  Hence  intense  heat  is  developed  along  the  carbon 
core.  Large  electric  furnaces  constructed  on  this  type  are 
now  in  practical  operation.  And  since  electricity  is  obtained 
in  many  localities  by  operating  dynamos  by  water,  new  in- 
dustries requiring  intense  and  continuous  heat  have  quickly 
sprung  into  existence  and  older  ones  have  been  remodeled. 
Some  of  these  plants  are  located  at  Niagara  Falls,  which  fur- 
nishes enormous  power  at  a  relatively  small  expense.  Several 


HEAT  175 

commercial  substances,  more  or  less  familiar,  are  manu- 
factured in  the  electric  furnace;  e.g.  calcium  carbide, 
carborundum,  alundum,  phosphorus,  carbon  disulphide, 
graphite,  and  silicon.  These  are  discussed  in  appropriate 
places. 

Measurement  of  Heat  Energy.  —  Every  substance  possesses 
a  certain  amount  of  chemical  energy,  but  there  is  no  way  of 
determining  the  total  amount.  We  can,  however,  measure 
that  part  of  the  chemical  energy  which  is  transformed  into 
heat  when  a  substance  or  set  of  substances  undergoes  a 
chemical  change.  Thus,  when  hydrogen  and  oxygen  com- 
bine chemically,  the  total  amount  of  chemical  energy  in  the 
two  gases  is  divided,  part  being  liberated  as  heat,  and  part 
being  locked  up  as  chemical  energy  in  the  water  formed ; 
and  the  liberated  heat  may  be  taken  as  a  measure  of  the 
chemical  energy  transformed  during  the  chemical  change. 
Heat  measurements  are  made  in  a  calorimeter.  This  appa- 
ratus consists  essentially  of  two  parts,  a  small  vessel  in  which 
the  substance  is  chemically  changed  and  a  larger  one  contain- 
ing water  in  which  the  small  vessel  is  immersed.  The  heat 
involved  in  the  chemical  reaction  changes  the  temperature  of 
the  water.  The  fuel  (i.e.  heat)  value  of  food,  coal,  etc.,  is 
found  by  means  of  a  calorimeter.  Heat  is  measured  in  cal- 
ories ;  a  calorie  (cal.)  is  the  quantity  of  heat  necessary  to 
raise  1  gm.  of  water  1°  C.  in  temperature  (15°-16°  C.  being 
the  degree  usually  taken).  For  example,  the  heat  liberated 
by  the  burning  of  1  gm.  of  hydrogen  is  34,200  cal.  Atten- 
tion has  already  been  called  to  the  high  temperature  of  the 
hydrogen  flame.  (See  Chapter  III.) 

Ordinary  chemical  equations  do  not  express  changes  in 
energy.  To  represent  heat  changes,  the  number  of  calories 
of  heat  involved  is  placed  after  the  equation,  together  with 
the  proper  sign,  thus  :  — 


176  INORGANIC   CHEMISTRY 

2H       +     0     =    H20     +     68,400  cal. 

Hydrogen        Oxygen          Water 
2  16  18 

This  is  called  a  thermal  equation,  and  it  means  that  68,400 
cal.  of  heat  are  liberated  when  2  gm.  of  hydrogen  unite  with 
16  gm.  of  oxygen  to  form  18  gm.  of  water.  In  some  chemi- 
cal changes  heat  is  absorbed.  Thus,  when  carbon  unites 
with  sulphur  to  form  carbon  disulphide,  heat  is  absorbed. 
The  equation  expressing  this  fact  is  — 

C    +     S2       =         CS2       -      19,600  cal. 

Carbon      Sulphur  Carbon 

12  64  Disulphide 

76 

Heat  evolved  or  absorbed  in  the  formation  of  a  mole  of  a 
compound  is  called  its  heat  of  formation.  If  heat  is  liber- 
ated in  the  formation  of  a  compound,  the  heat  is  desig- 
nated positive  (+) ;  and  the  compound  is  termed  exothermic. 
Heat  of  formation  which  is  absorbed  is  designated  negative 
(— );  and  a  compound  having  a  negative  heat  of  forma- 
tion is  said  to  be  endothermic.  Exothermic  compounds  are 
relatively  stable ;  they  can  be  decomposed  by  the  addition  of 
the  same  quantity  of  heat  liberated  by  their  formation. 
Thus,  68,400  cal.  of  heat  or  an  equivalent  quantity  of  energy 
must  be  added  to  18  gm.  of  water  to  decompose  it  into  2  gm. 
of  hydrogen  and  16  gm.  of  oxygen.  Such  heat  is  called  heat 
of  decomposition.  On  the  other  hand,  endothermic  com- 
pounds are  unstable  and  often  explosive.  They  decompose 
easily  with  the  liberation  of  heat.  The  heat  evolved  when 
organic  substances  are  ultimately  oxidized  to  carbon  dioxide 
and  water  is  called  heat  of  combustion.  Our  knowledge  of 
the  calorific  value  (page  269)  of  fuels  and  foods  is  based 
largely  on  measurements  of  their  heats  of  combustion.  In 
the  preceding  chapter  attention  was  called  to  the  fact  that 
the  interaction  of  equivalent  quantities  of  strong  acids  and 


ELECTRICITY  177 

bases  liberates  the  same  quantity  of  heat ;  that  is,  the  heat 
of  neutralization  liberated  by  the  union  of  hydrogen  and 
hydroxyl  ions  is  13,700  cal.  under  normal  conditions.  The 
thermal  equation  for  the  neutralization  of  sodium  hydroxide 
by  hydrochloric  acid  is  — 

NaOH  +  HC1  =  NaCl  +  H2O '+  13,700  cal. 

Applications  and  extensions  of  thermochemistry,  as  this 
branch  of  the  science  is  called,  will  be  made  in  subsequent 
chapters. 

ELECTRICITY  AND  CHEMICAL  CHANGE 

The  Relation  between  Electricity  and  Chemical  Action  has 
always  been  a  fascinating  subject.  Volta  constructed  his 
voltaic  pile  about  1800.  This  was  one  of  the  first  (perhaps 
the  first)  sources  of  an  electric  current.  In  May,  1800, 
Nicholson  and  Carlisle  decomposed  water  into  hydrogen 
and  oxygen  by  an  electric  current  obtained  from  a  thermo- 
pile. In  the  same  year  Cruikshank  obtained  lead  and 
copper  from  solutions  of  their  salts.  And  in  1807  Davy 
isolated  two  elements,  sodium  and  potassium,  by  passing 
an  electric  current  (obtained  from  a  large  battery)  through 
fused  sodium  hydroxide  and  potassium  hydroxide  respec- 
tively. From  that  time  until  the  present  day  the  relation 
between  electricity  and  chemical  change  has  engaged  the 
attention  of  chemists,  and  their  labors  have  built  up  a 
branch  of  chemistry  called  electrochemistry,  which  has 
recently  attained  considerable  commercial  importance. 

Transformations    of   Electrical    and    Chemical   Energy.  — 

Electricity,  like  heat  and  light,  is  readily  transformed  into 
chemical  energy,  and  vice  versa.     Certain  chemical  changes 
produce  electricity,  while  electricity  is  consumed  in  others. 
A  typical  illustration  of  the  transformation  of  chemical 


178  INORGANIC  CHEMISTRY 

energy  into  electricity  is  furnished  by  the 
voltaic  cell.     The  simplest  form  consists 
of  two  unlike  metals,  such  as  copper  and 
zinc,  connected    by    a    wire    and    partly 
immersed   in   a  vessel   containing   dilute 
sulphuric  acid  (Fig.  21).     When  the  con- 
FIG.  21.  — Simple  vol-  nected  metals  are  put  into  the  acid,  the 
taic  cell,    z  is  the  zjnc  slowiy  disappears  and  hydrogen  bub- 
zinc   and   C  is  the   ,  ,  ,,  „      . . 
copper                     bles  appear  on  the  copper.     Further  ex- 
amination shows  that  zinc  sulphate  is  also 
formed.     The  chemical  change  is  essentially  the  one  already 
described  under  hydrogen,  and  can  be  represented  thus  :  — 


Zn    +    H2SO4      =     2H     +     ZnSO4 

:>huric  Hydrogen  Zinc 

cid  Sulphate 


Zinc  Sulphuric  Hydrogen  Zinc 

Ac' " 


The  wire  becomes  electrified  and  exhibits  the  effects  of  an 
electric  current.  For  example,  it  becomes  warm  and  makes 
a  magnetic  needle  move.  The  source  of  the  electric  current 
is  the  chemical  action  between  the  acid  and  zinc.  The 
copper  is  necessary,  otherwise  the  product  of  the  chemical 
action  would  be  merely  heat.  Carbon  is  often  used  instead 
of  copper,  and  sulphuric  is  replaced  by  other  liquids,  such 
as  solutions  of  ammonium  chloride  or  potassium  hydroxide. 
The  liquid  chosen,  however,  must  be  one  that  will  interact 
with  zinc  or  its  substitute.  The  zinc  is  ultimately  trans- 
formed into  zinc  sulphate  or  some  other  chemical  compound, 
and  must  be  replenished;  the  solution  must  likewise  be 
renewed.  This  necessity  of  recharging  the  cell  is  a  striking 
proof  of  the  intimate  relation  between  chemical  change  and 
electricity.  Several  cells  joined  together  form  an  electric 
battery.  For  many  years  the  battery  was  the  chief  source 
of  the  electric  current ;  and  it  is  now  widely  used  to  generate 
the  currents  of  moderate  intensity  utilized  in  ringing  bells 


ELECTRICITY 


179 


and  operating  electrical  apparatus.  Powerful  currents,  as 
a  rule,  are  obtained  from  a  dynamo. 

The  transformation  of  electricity  into  chemical  energy  is 
illustrated  by  electrolysis.  In  Chapter  IX  considerable 
space  was  devoted  to  a  discussion  of  the  electrolysis  of  solu- 
tions; i.e.  to  the  chemical  changes  which  accompany  the 
passage  of  an  electric  current  through  a  solution.  But 
electrolysis  is  not  limited  to  solutions.  Fused  (i.e.  melted) 
substances  also  undergo  chemical  changes  when  subjected 
to  the  action  of  a  powerful  electric  current.  As  already 
stated,  electrolysis  is  accomplished  in  an  electrolytic  ceil. 
This  apparatus  differs  from  a  voltaic  cell  in  one  essential 
respect.  The  voltaic  cell  produces  electricity,  whereas  the 
electrolytic  cell  consumes  electricity.  Otherwise  both  have 
three  main  parts,  — the  containing 
vessel,  two  electrodes,  and  the 
electrolytic  solution  or  fused  elec- 
trolyte. A  simple  electrolytic  cell 
is  shown  in  Figure  22.  In  such  a 
cell  the  current  from  the  battery 
or  dynamo  enters  the  cell  through 
the  positive  (  +  )  electrode  or  anode  FIG.  2*2.  —  Simple  electrolytic 
and  leaves  through  the  negative  ^alde.tdTt  <£ 
(  — )  electrode  or  cathode.  Elec-  troiytic  solution, 
trodes  may  be  of  platinum,  copper, 

zinc,  mercury,  or  hardened  carbon ;  they  may  have  any 
shape, — rod,  wire,  sheet,  plate,  box,  crucible;  and  they 
also  may  be  solid,  liquid,  or  powder,  as  well  as  fixed  or 
movable.  The  electrodes  are  connected  by  wires  with  the 
source  of  the  electric  current,  and  serve  as  "doors"  — 
to  quote  Faraday  —  for  the  current  to  flow  into  and  out 
of  electrolytic  solution  or  the  electrolyte.  We  speak  of  a 
"current"  of  electricity  and  of  electricity  as  "flowing," 
although  we  do  not  know  the  fundamental  nature  of  elec- 


180  INORGANIC  CHEMISTRY 

tricity  nor  do  we  mean  really  that  it  flows,  like  a  river,  only 
in  one  direction.  The  anode  is  the  electrode  that  is  often 
consumed  or  worn  away,  either  mechanically  or  chemically, 
but  solids,  especially  metals,  are  often  deposited  upon  the 
cathode.  The  containing  vessel  likewise  may  have  any 
desired  shape,  and  is  usually  made  of  some  material  which 
will  resist  the  corrosive  action  of  chemicals;  e.g.  porcelain, 
slate,  or  soapstone.  Metallic  vessels  lined  with  various  sub- 
stances selected  to  withstand  intense  heat  are  also  used, 
since  fused  electrolytes  are  often  subjected  to  electrolysis. 

The  phenomena  exhibited  during  the  electrolysis  of  solu- 
tions have  been  described.  (See  Chapter  IX.)  Certain 
features  of  the  operation,  however,  should  be  recalled  at 
this  point,  since  they  apply  to  both  fused  and  dissolved 
electrolytes.  Employing  the  interpretation  offered  by  the 
theory  of  electrolytic  dissociation,  the  ions  of  the  electrolyte 
begin  to  migrate  as  soon  as  the  electric  current  enters  the 
cell;  the  electro-negative  anions  move  toward  the  electro- 
positive anode,  and  the  electro-positive  cations  move  toward 
the  electro-negative  cathode.  Reaching  their  respective 
electrodes,  the  ions  give  up  their  charges  and  become  atoms, 
atomic  groups,  or  molecules.  The  discharged  particles  may 
escape  as  gases,  dissolve  in  the  liquid,  or  attach  themselves 
to  the  electrodes;  very  often  secondary  chemical  changes 
occur  which  complicate  the  process  and  sometimes  cause 
serious  difficulties  in  industrial  applications  of  electrolysis. 

A  third  illustration  of  the  transformation  of  electrical 
energy  into  chemical  energy,  and  vice  versa,  is  furnished  by 
the  storage  cell.  This  cell  consists  essentially  of  two  grids 
of  lead  which  are  filled  with  a  mixture  of  sulphuric  acid  and 
lead  oxide  (PbO)  and  then  immersed  in  a  vessel  containing 
dilute  sulphuric  acid.  The  mixture  in  the  grids  soon  be- 
comes lead  sulphate  (PbSO4).  When  an  electric  current  is 
passed  through  the  cell  from  one  plate  to  the  other,  the 


ELECTRICITY  181 

hydrogen  ions  formed  by  the  ionization  of  the  sulphuric 
acid  migrate  to  the  cathode,  where  they  become  atomic 
hydrogen  and  by  their  chemical  interaction  with  the  lead 
sulphate  form  sulphuric  acid  and  metallic  lead  —  the  latter 
remaining  attached  to  the  grid  (cathode);  the  S04-ions 
migrate  to  the  anode,  where  they  become  ordinary  uncharged 
atomic  groups  and  by  their  interaction  with  the  lead  sul- 
phate form  sulphuric  acid  and  lead  dioxide  (PbO2)  — the 
latter  remaining  attached  to  the  grid  (anode).  These 
changes  continue  until  nearly  all  the  lead  sulphate  has 
been  altered  as  just  described.  If,  after  being  charged,  the 
grids  are  connected  by  a  wire,  a  current  of  electricity  will 
be  obtained  in  a  direction  opposite  to  that  used  in  charging 
the  cell,  and  the  chemical  changes  will  take  place  in  the 
reverse  direction.  The  cell  gradually  reverts  to  its  original 
condition,  and  must  be  recharged  if  a  current  is  again 
desired.  In  this  cell,  therefore,  there  is  a  complete  circuit 
of  transformation,  —  electricity  — >  chemical  energy  — >  elec- 
tricity. 

Industrial  Applications  of  Electrolysis.  —  The  earliest  in- 
dustrial application  of  electrolysis  was  in  electrotyping  and 
electroplating.  These  operations  consist  in  depositing  a  thin 
film  of  metal  upon  a  surface.  They  are  fundamentally  the 
same,  though  copper  is  the  only  metal  used  for  producing 
electrotypes.  Electrotypes  are  exact  reproductions  of  the 
original  object.  The  process  of  electrotyping  is  substantially 
as  follows  :  The  page  of  type,  for  example,  is  first  repro- 
duced in  wax.  This  exact  impression  is  next  covered  with 
powdered  graphite  to  make  it  conduct  electricity.  The 
coated  mold  is  then  suspended  as  the  cathode  in  an  acid 
solution  of  copper  sulphate;  the  anode  is  a  plate  or  bar  of 
copper.  When  the  current  is  passed  through  the  system, 
electrolysis  occurs;  copper  is  dissolved  from  the  anode  and 


182  INORGANIC  CHEMISTRY 

deposited  on  the  mold  in  a  film  of  any  desired  thickness. 
The  exact  copper  copy  is  stripped  from  the  mold,  backed 
with  metal,  and  used  instead  of  the  type  itself.  By  this 
process,  exact  copies  of  expensive  wood  engravings  can  be 
cheaply  reproduced  and  type  can  be  saved  from  the  wear 
and  tear  of  printing.  Most  books,  magazines,  and  news- 
papers are  now  printed  from  electrotypes.  The  process  of 
electroplating  differs  from  electrotyping  in  only  one  essen- 
tial; viz.  in  electroplating  the  deposited  film  is  not  removed 
from  the  object.  The  object  to  be  plated  is  carefully  cleaned 
and  made  the  cathode;  the  anode  is  a  bar  or  plate  of  the 
metal  to  be  deposited.  When  the  current  passes  through 
the  system,  the  metal  is  firmly  deposited  on  the  object. 
The  electrolysis  would  take  place,  of  course,  if  any  anode 
were  present;  but  anodes  of  the  metal  to  be  deposited  are 
usually  used  to  prevent  the  solution  or  "bath"  from  weaken- 
ing. They  accomplish  the  purpose  by  replenishing  the  solu- 
tion with  metal  as  fast  as  it  is  removed  and  deposited  upon 
the  cathode.  Silver,  nickel,  and  gold  are  the  usual  metals 
used  in  electroplating.  (See  these  metals.) 

Electroplating  and  electrotyping  have  been  done  since 
about  1840.  It  is  only  within  the  last  ten  or  fifteen  years, 
however,  that  the  electric  current  has  been  profitably  ap- 
plied in  many  industries.  But  during  this  time  the  develop- 
ment of  electrochemistry  has  been  very  marked.  The  largest 
of  these  industries  is  the  refining  of  copper.  The  process  is 
similar  to  that  described  under  electrotyping.  Other  metals, 
such  as  gold,  silver,  and  lead,  are  extracted  from  their  ores 
and  purified  by  electricity,  though  the  older  processes  are 
still  used.  The  aluminium,  magnesium,  and  sodium  of 
commerce  are  now  manufactured  by  passing  an  electric 
current  through  their  fused  compounds.  Chlorine,  potas- 
sium chlorate,  potassium  hydroxide,  and  sodium  hydroxide 
are  some  of  the  other  important  industrial  products  of 


ELECTRICITY  183 

electrolysis.     These  electrochemical'  processes  and  products 
will  be  fully  discussed  in  the  appropriate  places. 

Measurement  of  Electrical  Energy.  —  Faraday  was  the 
first  scientist  to  make  a  thorough  study  of  electrolysis.  He 
found  that  a  given  current  of  electricity  liberated  different 
but  definite  amounts  of  the  chemical  elements.  Thus, 
the  current  which  liberates  1  gm.  of  hydrogen  also  liber- 
ates 8  gm.  of  oxygen,  35.46  gm.  of  chlorine,  107.88  gm.  of 
silver,  31.78  gm.  of  copper,  and  so  on.  These  numbers  are 
identical  with  the  chemical  equivalents  of  these  elements. 
(Compare  Equivalents,  Chapter  XIV.)  Faraday  called  them 
electrochemical  equivalents,  to  emphasize  their  chemical 
and  electrical  relationship.  •  But  the  term  electrochemical 
equivalent  now  means  the  weight  of  an  element  deposited 
or  liberated  by  a  specified  current  in  a  certain  time  (1  ampere 
in  1  second).  For  example,  the  electrochemical  equivalent 
of  hydrogen  is  .000010441  gm.,  of  oxygen  is  .00008287 
and  sometimes  .00016574,  of  copper  is  .0003294  and  some- 
times 0.0006588,  of  silver  is  .001118.  This  general  relation 
is  often  stated  as  Faraday's  law,  thus :  — 

When  the  same  current  of  electricity  is  passed  through  solu- 
tions of  different  electrolytes,  the  ratio  of  the  quantities  of 
liberated  products  is  the  same  as  that  of  their  chemical  equiva- 
lents. 

Faraday  also  showed  that  — 

The  amount  of  decomposition  —  the  chemical  work,  we 
might  say  —  is  proportional  to  the  total  amount  of  electricity 
used.  It  makes  no  difference  whether  the  current  is  strong 
or  weak,  nor  whether  the  time  of  its  flow  is  long  or  short. 
A  certain  quantity  of  electricity  will  do  so  much  chemical 
work  —  no  more  and  no  less.  Thus,  a  given  quantity  of 
electricity  passed  through  copper  sulphate  solution  always 
deposits  the  same  weight  of  copper  at  the  cathode. 


184  INORGANIC  CHEMISTRY 

These  two  principles  of  Faraday  are  at  the  foundation 
of  all  electrochemical  industries.  Their  importance  can 
hardly  be  overestimated. 

OTHER  CONDITIONS  AFFECTING  CHEMICAL  ACTION 

Substances  vary  in  their  tendency  to  undergo  chemical 
change.  Some,  like  oxygen,  are  active  elements;  i.e.  they 
unite  directly  with  other  substances,  liberate  energy  rapidly, 
and  form  comparatively  stable  compounds.  Others,  like 
nitrogen,  are  inert,  and  form  somewhat  unstable  compounds 
by  indirect  processes.  Not  only  does  the  chemical  activity 
of  substances  differ,  but  the  same  substance  is  not  equally 
active  toward  all  others.  Thus,  iron  unites  readily  with 
sulphur  and  bromine,  but  not  with  copper  or  mercury. 
Nevertheless,  certain  substances  which  are  inert  under  some 
conditions  may  become  active  under  special  conditions. 
That  is,  chemical  action  depends  not  only  upon  the  specific 
attraction  of  interacting  substances,  i.e.  their  chemical 
affinity  as  it  is  sometimes  called,  but  also  upon  the  special 
conditions  under  which  the  reaction  occurs.  Chemical 
changes  which  take  place  very  slowly,  as  we  have  already 
seen,  sometimes  proceed  with  astonishing  rapidity  under  the 
influence  of  heat,  light,  or  electricity ;  that  is,  the  velocity  of 
the  reaction  is  increased.  In  many  chemical  changes,  how- 
ever, the  velocity  of  the  reaction  is  not  influenced  solely  by 
one  of  the  familiar  forms  of  energy,  but  also  by  other  factors, 
e.g.  solution,  concentration,  catalysis,  and  equilibrium.  By 
the  velocity  of  a  reaction  we  mean  the  amount  of  substance 
that  reacts  in  a  given  time. 

The  Effect  of  Solution  on  chemical  change  was  discussed 
in  Chapter  IX.  Many  chemical  changes  are  carried  out  in 
aqueous  solutions  because  dissolved  substances  are  in  a  con- 
dition especially  favorable  for  interaction.  Therefore  "  wet 


CATALYSIS 


185 


processes,"  as  they  are  sometimes  called,  are  exceedingly 
important  in  chemical  analysis.  The  velocity  with  which  a 
reaction  proceeds  is  often  greatly  influenced  by  concentration  ; 
that  is,  by  the  quantity  of  substance  in  a  given  volume.  If 
the  concentration  is  increased,  the  reaction  takes  place  more 
rapidly;  that  is,  in  general,  the  greater  the  concentration, 
the  greater  the  reaction  velocity.  A  concrete  illustration 
will  make  this  point  clearer.  An  acid  solution  of  potassium 
iodide  and  starch  turns  blue  when  an  oxidizing  agent  is 
added,  owing  to  the  interaction  of  the  starch  and  the  liber- 
ated iodine.  If  the  potassium  iodide  solution  and  oxidizing 
solution  are  added  to  the  same  volume  of  water,  the  time 
required  to  produce  a  standard  blue  color  will  be  a  relative 
measure  of  the  reaction  velocity.  The  following  table  shows 
the  result  of  an  experiment :  — 

MEASUREMENT  OF  REACTION  VELOCITY 


SOLUTION 

OXIDIZING  SOLUTION 

POTASSIUM  IODIDE 
SOLUTION  (cc.) 

TIME  ix  MINUTES 

I 

5 

5 

4 

II 

10 

5 

2 

III 

5 

10 

2 

IV 

10 

10 

1 

From  the  table  it  is  evident  that  as  the  concentration  of 
solution  IV  is  four  times  that  of  solution  I,  the  time  for  IV 
is  one  fourth  that  of  I ;  solutions  II  and  III  have  the  same 
concentration  and  their  times  are  the  same. 

Catalysis.  —  The  velocity  of  some  reactions  can  be  altered 
by  the  presence  of  certain  substances  that  apparently  do 
not  participate  in  the  chemical  change  and  can  be  recovered 
unchanged  after  the  action  has  ceased.  Such .  substances 
are  called  catalytic  agents  or  catalyzers,  and  their  effect 


186  INORGANIC  CHEMISTRY 

upon  chemical  action  is  called  catalysis  or  catalytic  action. 
For  instance,  potassium  chlorate  yields  oxygen  slowly  when 
heated  to  about  350°  C.,  but  if  powdered  manganese  dioxide 
is  added,  the  gas  is  evolved  rapidly.  Furthermore,  the 
manganese  dioxide  can  be  recovered  after  the  experiment 
merely  by  dissolving  out  the  residual  potassium  chloride. 

Similarly,  hydrogen  dioxide  (H2O2)  decomposes  very 
slowly  in  the  air,  but  if  manganese  dioxide  is  added,  the 
decomposition  proceeds  rapidly  and  bubbles  of  oxygen  gas 
can  be  seen  rising  through  the  liquid.  Likewise,  a  mixture  of 
hydrogen  and  oxygen  at  the  ordinary  temperature  reveals  no 
tendency  toward  chemical  action,  but  if  a  little  powdered 
platinum  is  added,  the  gases  combine  with  almost  explosive 
violence.  Again,  sulphur  dioxide  gas  (S02)  and  oxygen 
united  very  slowly,  but  when  a  purified  and  properly  cooled 
mixture  of  these  gases  is  passed  over  finely  divided  platinum, 
the  gases  unite  rapidly  and  form  sulphur  trioxide  (SO3). 
This  rapid  transformation  is  utilized  in  one  process  of  manu- 
facturing sulphuric  acid.  Water  vapor  in  traces  is  regarded 
by  some  authorities  as  a  catalytic  agent.  Thus,  many  gases, 
especially  hydrogen  chloride  (HC1)  and  ammonia  (NH3),  do 
not  unite  when  perfectly  dry,  but  if  a  trace  of  water  is  added, 
the  reaction  proceeds  as  ordinarily  observed ;  i.e.  white 
fumes  of  ammonium  chloride  (NH4C1)  are  formed.  In  many 
technical  processes  reactions  are  hastened  by  catalyzers, 
often  in  relatively  small  quantities,  e.g.  mercury,  nickel, 
chlorides  of  copper,  zinc,  and  aluminium,  dilute  acids,  and 
enzymes.  Sometimes  a  catalyzer  retards  a  reaction ;  such 
substances  are  called  negative  catalyzers.  Catalysis  is  an 
important  phenomenon. 

Reversible  Reactions  and  Chemical  Equilibrium.  —  Hitherto 
we  have  described  chemical  reactions  as  if  they  proceeded 
to  completion.  Indeed,  they  do  apparently,  for  in  many 


CHEMICAL  EQUILIBRIUM  187 

experiments  conditions  are  chosen  which  necessitate  action 
until  one  or  more  of  the  interacting  substances  is  exhausted. 
As  a  matter  of  fact  many  reactions  are  reversible;  that  is, 
they  proceed  in  one  direction  under  one  set  of  conditions 
and  in  the  opposite  direction  under  another  set  of  condi- 
tions. For  example,  in  the  manufacture  of  oxygen  from 
barium  oxide  a  reversible  reaction  occurs.  When  barium 
oxide  is  heated  in  the  air  to  about  700°  C.,  the  chemical 
change  is  represented  thus  :  — 

BaO     +     O     =     Ba02 

Barium          Oxygen          Barium 
Oxide  Dioxide 

If  the  air  supply  is  cut  off  and  the  pressure  (in  the  retorts) 
is  reduced,  the  chemical  change  is  reversed  and  may  be 
represented  thus :  — 

BaO2     =     O     +     BaO 

Barium          Oxygen          Barium 
Dioxide  Oxide 

Another  illustration  is  provided  by  Lavoisier's  famous  ex- 
periment described  under  Hydrogen  (see  Chapter  III).  He 
passed  steam  over  red-hot  iron  and  obtained  hydrogen  and 
iron  oxide;  the  equation  for  the  chemical  change  is  — 

4H2O     +     3Fe     =     8.H     +     Fe3O4 

Water  Iron  Hydrogen  Iron 

Oxide 

When  hydrogen  is  passed  over  hot  iron  oxide,  the  chemical 
change  is  reversed,  thus  :  — 

8H     +     Fe304     =     3  Fe     -f-     4H20 

Hydrogen  Iron  Iron  Water 

Oxide 

We  might  conclude  that  in  reversible  reactions  the  chemi- 
cal change  can  proceed  in  either  direction  to  completion.  As 
a  matter  of  fact  the  actual  or  effective  chemical  change  de- 
pends upon  the  conditions.  Thus,  in  the  interaction  of 


188  INORGANIC  CHEMISTRY 

steam  and  iron,  if  the  hydrogen  is  continuously  swept  out 
of  the  tube  by  the  steam,  no  reduction  of  the  iron  oxide 
occurs ;  or  if  in  the  interaction  of  hydrogen  and  iron  oxide 
the  steam  is  continuously  removed,  no  oxidation  of  iron  will 
occur.  But  when  the  experiment  is  performed  in  a  closed 
tube,  the  result  is  an  equilibrium  between  the  two  oppos- 
ing reactions.  Not  only  does  the  interaction  of  the  steam 
and  iron  produce  hydrogen  and  iron  oxide,  but  conversely, 
the  interaction  of  the  hydrogen  and  iron  oxide  forms  steam 
and  iron.  The  reactions  proceed  simultaneously  in  the  same 
tube  until  chemical  equilibrium  is  reached ;  that  is,  not  to 
completion  in  either  direction,  but  to  such  a  point  in  both 
directions  that  there  is  no  further  alteration  in  the  weights 
of  the  substances  actually  participating  in  the  change.  Let 
us  illustrate  more  definitely.  If  iron  and  steam  are  heated 
in  a  closed  tube,  the  velocity  of  the  forward  reaction  (i.e.  the 
transformation  of  steam  and  iron  into  iron  oxide  and  hydro- 
gen) gradually  diminishes  until  it  apparently  stops,  although 
some  material  is  still  available  for  chemical  action ;  con- 
versely, the  velocity  of  the  reverse  reaction  (i.e.  the  reforma- 
tion of  steam  and  iron  from  iron  oxide  and  hydrogen)  gradu- 
ally increases  until  it  likewise  apparently  stops,  although 
additional  material  is  available.  Continued  heating  pro- 
duces no  change  in  the  relative  weights  of  the  four  sub- 
stances in  the  tube;  that  is,  equilibrium  has  been  reached. 
This  means  that  the  velocities  of  the  opposing  reactions  are 
equal.  In  other  words,  there  is  no  further  accumulation  of 
any  of  the  reacting  substances,  because  the  opposing  re- 
actions are  proceeding  at  the  same  rate ;  both  the  forward 
and  reverse  reactions  are  still  taking  place,  but  one  undoes 
the  work  of  the  other,  so  to  speak.  Chemical  equilibrium, 
then,  is  that  state  reached  in  a  reversible  reaction  when  the 
further  accumulation  of  any  of  the  reacting  substances  is  pre- 
vented by  the  equal  velocities  of  the  two  opposing  reactions. 


CHEMICAL  EQUILIBRIUM  189 

In  equations    representing    reversible  reactions   oppositely 
directed  arrows  (read  "  equals  reversibly  ")  are  used,  thus  :  — 

H2  +  I2  ^±  2  HI. 

The  velocity  of  a  reaction  is  affected  by  several  factors, 
especially  the  concentration  of  the  reacting  substances. 
The  law  covering  the  relation  between  the  concentration 
and  the  velocity  of  a  reaction  is  often  called  the  law  of  mass 
action,  and  may  be  stated  thus : 

At  a  constant  temperature  the  velocity  of  a  reaction  is  pro- 
portional to  the  molecular  concentration  of  each  reacting 
substance. 

By  concentration  of  the  reacting  substance  is  meant  that 
relative  portion  of  each  substance  actually  available  for 
chemical  action  at  a  given  time.  By  molecular  concentra- 
tion is  meant  the  number  of  moles  (i.e.  gram-molecular 
weights)  in  a  liter.  The  law  finds  its  best  application  in  the 
case  of  gases  and  solutions  because  these  systems  of  sub- 
stances, being  homogeneous  mixtures,  are  in  a  favorable 
condition  for  chemical  action ;  moreover,  the  concentration 
of  each  ingredient  of  a  homogeneous  mixture  can  readily  be 
expressed.  In  applying  the  law  to  heterogeneous  mixtures, 
e.g.  steam-iron-iron  oxide-hydrogen,  it  is  customary  to  con- 
sider only  the  gases  because  the  concentration  of  the  par- 
ticipating portion  of  each  solid  —  the  active  mass,  as  it  is 
sometimes  called  —  is  practically  constant. 

Let  us  apply  the  law  of  mass  action  to  a  gaseous  reaction, 
viz.  the  combining  of  hydrogen  and  iodine  to  form  hydriodic 
acid.  The  equation  correctly  wrftten  for  our  present  pur- 
pose is :  — 

H2  +  I2  =  2  HI. 

Let  (a)  and  (6.)  represent  the  molecular  concentration  of  the 
hydrogen  and  iodine  respectively.     The  velocity  of  the  re- 


I'M)  i    rmci      K    •  in  .11   1  i:  . 


.i'h'.  l,    i       proportional     In    •  -:i«-|i    li.nl.-riil;,,     ,  .  ,i  ,,  ,  ,  ,  I  ,  .,  I  ,,  ,,,    ;i,,,| 

It"   "Inn       h,     Ilinr     piodllfl.        '|'|,,.     V»'lor|ly     alno    I|C|)(-||(|H    OH 

i|'  1  1  I  i'  loI'M  aM  the  .  •  peril,,-  ;,  |||  mil'-.    .,|   I  IK    i  ,.,,•!  n,g  HI|l)Hl.ii.n>  ' 

h  nip.  i.ii  UK  ,  .ui'l  •  .ii.il  ,  i  ic  agents.  Al  a  Kiven  temperature 
Mi-  effect  "I  tin  rninpli-x  Mel,  o  I  fuetoi  i  rojiHln.nl,  ;md  tin- 
.11..  i  i  ..ii.n  r;,||.-,|  i|,,.  affinity  constant.  Tin- 
coriMtant  IIIIH  :>  ililTcrcn!  vulur  l«>i  <  ...-I,  n-action  .-iinl 

I;:     l«-|,|c.  •.rilled     l.y     /•  ,     /.'|,     ''I,-          llrnrr     llir    r«  j  I  j;,  I  |,  ,,  ,     |,,,-     ||,r 
'il    ,       '-I      Ilir      H'.'.cllOI.      |,rl  WCCII      ll\.|l«  ,.-.•,,     „,,„!      ,,,,||IH.     |,|!- 


(1)  volooity  -  (a)  X  (6)  X  k. 

l.rl     H       ii'    .1     COIll  ulri     Ih,      r,      ,   ,     ,      rrnrlion  !,.• 

|)ONil.ioit  of  hydriodic  :i,-i,|   mh.  hydrogen  :md   iodine, 

,\   Hi, 


2  III        II.,    I    I,. 

In  •  •.pri'MHiliK  I'M-  \cl.M-ilv  of  MM-  <!«•(•(,  inposilion  (,f  liydiindir 
:i.'|.|,  two  fjlH,M  HillHt  l»<  l,|.|,  it,  h,  ;,,-,-,  ,i|ii  I  ;  (I)  luo 
l.|..|rri||r  ::  <,|  1  1  \  <  1  1  H  ..  I  |r  ;i  n,  I  C'  III)  ;i|,  |,,n,,,,|  from  U  HJII^Ic 

in.,1.  ,  -ulr  each  of  hydror.'-M  (II.)  ;in.|  iodine  (li),  :md  C2)  MM- 
iillinil  y  roil  l:ilil  ItHH  :i  dlfh-irnl  Villlir  froill  lli:il  111  '-'I  III  ll.r 

forward    reaction       ('on  (•(|ii(*nl.ly,   if   iln-   molrr-iil;i.r  rom-m 
h.  in..  n    ni    hydriodid    n.  -id    i     reprMntad    l»y   (c),   HH-II   UK- 

\-r|nr|lv  "I  Hl<  MMrlloll  IM  pI'OpOI'l  JOM.'d  In  (<')"  :illd  h)  /,',; 
;i  li<  I  I  lir  ri  |l|;i  I  inn  I  irr«  due 

(2)   vrlnrily    -   (c)8  X   k,. 

Sinn-  .ii  «  •«  1  1  n  1  1  1.  1  1  inn  ilm  rclociticHof  UK    forward  :md  rovcn  <• 
an-  r<|iml,  Mn-n  : 


(«)  X  (b)  Xk  -  (c)«XK,. 
Therefore 

(n)  X  (M        /,',         ... 


CHEMICAL  r.QUUBRIUM  101 

This  new  constant  K  is  called  tin-  equilibrium  constant. 
Sni'-c  it,  is  UK-  ratio  of  the  i,wo  affinity  corn  tanl  ,  it  nuin'-n 
ca,l  valu';  JH  conntant  at  a  constant  temper  a.l  ure.  That  IH, 
whatever  the,  original  concent  .rations  of  the  reacting  Huh- 
stances,  reactions  will  occur  such  thai  ;.t  equilibrium  the 
concentrations  will  give  the  a.mc  va.luc  for  t.h*-  con  i.-mt.  K. 

riirnniral  ^<juilihrium  IH  a  0eriHJtivr:  relation  [>etwe,en 
weigfits  of  Hijh«tance.-:  a./jf|  it  in  oanily  influeficed  hy  a.  cfiangf; 
in  the  conditions.  In  otfjr-r  wonlx,  a  change  in  f  <  i/,p'  r  ;»,ture 
or  r:oncenir;i.tion  favors  one  <>!'  \\\c1  opposing  reactions  and 
thereby  profluee»  anothfrr  state  <>\'  erjijilihriurri  wliich  corrrr- 
Hi>onHs  to  I  he  new  conditions.  The  cha.nge,  in  equilibrium 
ca.u~er|  hy  a.  cha.ngo  in  conditions  is  called  displacement  of 
equilibrium.  The  two  main  factors  that  cauftf;  dinplaccment 
cjf  equilibrium  are  temperature  and  concentration.  A  riso 
in  temperature  increases  the  velocity  of  a  reaction  ;  the 
velocity  in  many  eawn  in  doubled  hy  a  HMO  of  ten  dcgrwH. 
An  a  rule,  however,  the  velocities  of  the  opposing  reactions 
in  a  reversible  rcactio  ffected  quite  differently  hy  a 

change  in  temperature.     H'nce  efjuilibriurn  is  usually  dis- 
placed  hy  a  change  in   temperature. 

Th<-  direction  i/i  which  equilibrium  is  displaced  hy  a 
change  in  tanperfttttTC  <\<-\><-n<\  <>\\  which  reaction  —  the 
forward  or  the.  reverse  absorbs  heat.  The  reaction  that 
absorbs  heat  is  favored  hy  rise  of  temperature,  arid  vice 
versa;  thun  in  the  case  of  the  revcrsib  --  , 

4  11,0  +3Fel8  ff 


the  reverne  reaction  ^ifidic;j.ted  by  the  lower  arrow;  in  favored 
because  it  absorb*  heat.     The  rr-|;>.iion  of  temperature  and 
di.Mpla.cer/M-nt.   of   equilibrium    is  a  special    caw;  of  a   b 
gerM-ra.lixation   known  as   Le  Chatelier's  Law  Cor  theorem;, 
which  may  be    iat«-d  thus  : 

When  a  nyntem  in  &juilibrium  itt  nubjecl«l  i<>  n  f-hfinge  of 


192  INORGANIC  CHEMISTRY 

conditions  (such  as  temperature  and  pressure)  the  system  alters 
in  the  way  that  neutralizes,  or  tends  to  neutralize,  the  effect  of 
the  change. 

The  influence  of  a  change  in  concentration  on  equilibrium 
is  apparent  from  a  consideration  of  a  concentration  equation, 
e.g.  the  equation  expressing  equilibrium  in  the  reversible 
reaction  involving  hydrogen,  iodine,  and  hydriodic  acid  :  - 

(o)  X  (6)  =  K 

(c)2 

Suppose  the  concentration  represented  by  (a)  is  changed, 
then  the  numerical  value  of  the  numerator  will  be  changed 
and  reactions  must  take  place  until  the  concentrations  are 
such  that  the  new  state  of  equilibrium  will  give  the  same 
value  of  the  constant  K.  In  other  words,  equilibrium  is 
displaced  by  changing  concentration.  Change  in  concen- 
tration may  be  brought  about  in  several  ways.  One  way 
is  the  removal  of  one  product  of  the  reaction.  (1)  If  the 
iodine  from  the  decomposition  of  hydriodic  acid  is  partly 
concentrated  by  cooling  one  end  of  the  tube,  this  iodine  no 
longer  participates  in  the  reaction,  and  the  concentration 
of  the  active  iodine  will  be  diminished.  That  is,  equilibrium 
is  displaced  and  the  decomposition  of  hydriodic  acid  must  go 
much  further  before  a  new  state  of  equilibrium  is  estab- 
lished. (2)  In  the  experiment  with  iron  and  steam,  if  the  tube 
is  opened  and  a  current  of  steam  is  introduced,  the  hydrogen 
which  is  formed  by  the  forward  reaction  will  be  removed  and 
the  reaction  will  continue  until  the  iron  is  oxidized.  That 
is,  the  removal  of  the  hydrogen  displaces  the  equilibrium 
and  thereby  permits  the  forward  reaction  to  proceed  to  com- 
pletion. Similarly,  if  hydrogen  is  introduced  into  the  tube, 
the  steam  will  be  removed  and  the  reduction  of  the  iron  by 
the  hydrogen  proceeds  to  completion.  In  other  words,  the 
entire  removal  of  one  product  of  the  reaction  displaces  the 


CHEMICAL  EQUILIBRIUM  193 

equilibrium  and  allows  the  reaction  to  proceed  to  completion. 
(3)  Removal  of  one  product  of  a  reaction  is  readily  accom- 
plished by  establishing  certain  reactions  in  solutions.  If 
one  of  the  products  is  insoluble,  it  is  removed  from  the 
sphere  of  action.  When  silver  nitrate  and  hydrochloric 
acid  are  dissolved  in  water,  the  solution  contains  at  first 
the  ions  Ag+,  NO3~,  H+,  and  Cl~,  and  undissociated  mole- 
cules of  silver  nitrate  (AgNO3)  and  hydrochloric  acid  (HC1). 
But  the  ions  Ag+  and  Cl~  at  once  form  molecules  of  insoluble 
silver  chloride  (AgCl),  thereby  displacing  the  equilibrium 
between  the  ions  and  their  undissociated  molecules.  The 
removal  of  these  ions  permits  the  dissociation  of  more  and 
more  silver  nitrate  and  hydrochloric  acid  molecules,  and  the 
precipitation  of  silver  chloride  continues  until  the  silver  or 
chlorine  ions  are  practically  exhausted.  (4)  A  gaseous  or 
readily  volatilized  product  of  a  reaction  can  be  removed  by 
raising  the  temperature ;  two  typical  examples  are  the  prep- 
aration of  hydrogen  chloride  and  nitric  acid  (see  pages 
204,  217). 

Another  way  to  change  concentration  and  thereby  dis- 
place equilibrium  is  actually  to  change  the  quantity  of 
(1)  the  solvent  or  of  (2)  an  ionic  substance  in  a  given  solu- 
tion. Consider  a  sodium  chloride  solution.  The  solution 
contains  molecules  (NaCl)  and  ions  (Na+  and  Cl~),  and 
the  molecules  are  in  equilibrium  with  the  ions,  thus :  — 

NaCl  ^±  Na+  +  Cl~. 

The  degree  of  dissociation  of  the  molecules  depends  on  the 
concentration  of  the  solution  (and  the  temperature).  The 
molecular  concentration  of  the  sodium  chloride  may  be 
represented  by  C  and  the  ionic  concentration  by  Ci  (for 
Na+)  and  C2  (for  Cl~) .  Then  we  may  write  :  — 

Ci  X  C2  =K 


194  INORGANIC  CHEMISTRY 

(1)  Now  if  water  is  added  to  the  solution,  some  of  the  mole- 
cules will  dissociate  into  ions,  and  the  molecular  and  ionic 
concentrations  will  diminish  as  dilution  proceeds,  though  the 
ionic  concentration  will  diminish  less  rapidly  than  the  molec- 
ular. On  the  other  hand,  if  water  is  removed  from  the 
solution  by  evaporation,  the  concentrations  will  increase  — 
the  molecular  concentration  in  this  case  increasing  more 
rapidly  than  the  ionic.  Equilibrium  is  displaced  in  both 
operations  —  diluting  and  concentrating  —  by  the  changes 
in  concentration,  because  the  reactions  take  place  until  the 
concentration  fraction  has  the  same  value  for  K.1  (2)  Sup- 
pose hydrogen  chloride  (or  concentrated  hydrochloric  acid) 
is  added  to  a  solution  of  sodium  chloride.  The  latter  solu- 
tion originally  contained  sodium  chloride  molecules  (NaCl), 
sodium  ions  (Na+),  and  chloride  ions  (Cl+).  The  hydrogen 
chloride  (or  hydrochloric  acid)  provides  hydrogen  chloride 
molecules  (HC1),  hydrogen  ions  (H+),  and  chloride  ions 
(Cl~).  Some  of  the  sodium  ions  must  unite  with  chloride 
ions  to  form  sodium  chloride  molecules  in  order  to  maintain 
equilibrium  in  the  solution.  Now  if  the  molecular  concen- 
tration of  the  sodium  chloride  was  very  large  in  the  first 
solution,  the  newly  formed  sodium  chloride  molecules  will 
be  in  excess  of  the  amount  that  can  dissolve  in  the  water, 
and  hence  some  sodium  chloride  must  be  precipitated. 

A  method  of  displacing  equilibrium  similar  to  (2)  above 
is  quite  effective  in  the  case  of  substances  which  are  not 
very  soluble  in  water.  Such  substances  readily  form  satu- 
rated solutions  of  rather  small  molecular  concentration. 
Hence  the  equilibrium  between  the  molecules  and  ions  is 
readily  displaced.  Thus,  although  barium  sulphate  is  very 
slightly  soluble  in  water,  it  dissolves  to  some  extent ;  and 
the  equilibrium  equation  may  be  written  :  — 

1  The  value  of  K  is  not  constant  in  the  case  of  strong  electrolytes  because 
certain  factors,  not  yet  understood,  cannot  be  incorporated  in  the  equation. 


CHEMICAL  EQUILIBRIUM  195 


X  (S0 


(BaS04) 


Now  when  barium  chloride  solution  is  added  to  sodium  sul- 
phate solution,  barium  sulphate  is  precipitated  because  a 
saturated  solution  of  barium  sulphate  is  almost  instantly 
formed  and  all  additional  barium  sulphate  must  be  precipi- 
tated. This  precipitation  may  be  interpreted  in  another 
way.  In  a  saturated  solution  at  a  given  temperature  the 
concentration  of  the  solute  (in  this  case  barium  sulphate)  is 
constant  (Ki).  Hence  the  equation  may  be  written  :  — 

(Ba++)  X  (S04—  )  =  K  X  K!  =  K'. 

In  other  words,  the  product  of  the  ionic  concentrations  in  a 
saturated  solution  is  constant  (K').  This  new  constant 
(K')  is  called  the  solubility  product  of  barium  sulphate. 
When  the  product  of  the  ionic  concentrations  in  any  solu- 
tion exceeds  the  numerical  value  of  the  solubility  product, 
some  of  the  ions  unite  to  form  a  precipitate.  As  a  rule, 
when  solutions  of  two  electrolytes  are  mixed,  a  double  de- 
composition involving  precipitation  takes  place  if  the  product 
of  the  concentrations  of  any  two  ions  exceeds  the  solubility 
product  of  the  salt  formed  by  their  combination. 

PROBLEMS 

1.  Calculate  the  affinity  constant  of  an  acetic  acid  solution  hav- 
ing a  total  molar  concentration  of  .1  and  a  molar  concentration  of 
each  ion  of  .0013.  Ans.  .0000171. 

2.  The  molar  solubility  of  potassium  chlorate  at  18°  C.  is  .52 
and  the  per  cent  of  ionization  is  70  (i.e.  .70).     Calculate  the  solu- 
bility product.  Ans.  .13. 

3.  If  the  molar  solubility  of  barium  sulphate  is  .00001  and  the 
ionization  is  .98,  what  is  the  solubility  product? 

4.  Calculate  the  solubility  product  of  calcium  hydroxide  if  its 
molar  solubility  is  .02  and  its  ionization  is  .88. 


CHAPTER   XII 


Chlorine  and  Hydrochloric  Acid 

CHLOKINE  is  an  important  gaseous  element,  and  its  com- 
pounds are  useful,  especially  hydrochloric  acid,  sodium 
chloride,  and  bleaching  powder. 

Occurrence.  — Free  chlorine  is  never  found  in  nature, 
but  its  compounds  are  widely  distributed,  the  most  abun- 
dant being  sodium  chloride.  Many 
compounds  of  chlorine  with  potas- 
sium, magnesium,  and  calcium  are 
found  in  the  deposits  at  Stassfurt  in 
Germany.  (See  Potassium.)  About 
2  per  cent  of  the  total  amount  of 
matter  in  the  ocean  is  chlorine,  and 
the  salts  found  in  sea  water  con- 
tain about  55  per  cent  of  chlorine. 
Silver  chloride  —  "  horn  "  silver  - 
is  mined  as  an  ore  in  the  United 
States  and  Mexico. 

Preparation.  —  Chlorine    is    pre- 
pared in  the  laboratory  by  heating 

FIG.  23. -Apparatus  for  pre-   &    m[xiurQ    of     manganese    dioxide 
paring   chlorine.      The    gas          ,.',.-.        ,  ,      .  •  ,     /T?-        Oo\ 

is  generated  in   A,  passes  and   hydrochloric    acid 

through  CE  to  the  bottom  This  method  was  used  by  Scheele, 

of   the   bottle    O    and    dis-   who    discovered    the    gas   in    1774. 
places  the  air,  which  escapes 

through  a  hole  in  the  cover  The  equation  for  the  preparation 
f.  of  chlorine  is  — 

196 


CHLORINE   AND   HYDROCHLORIC  ACID          197 
Mn02      +      4HC1      =     2C1     +     MnCl2    +    2  H2O 

Manganese          Hydrochloric          Chlorine          Manganese  Water 

Dioxide  Acid  Bichloride 

This  is  an  oxidizing  process,  since  the  hydrogen  of  the 
hydrochloric  acid  is  oxidized  to  water,  thereby  liberating 
a  part  of  the  chlorine  of  the  acid  as  free  chlorine  gas. 
Other  oxidizing  substances  besides  manganese  dioxide  may 
be  used,  such  as  potassium  chlorate  (KC1O3),  potassium 
dichromate  (K2Cr207),  red  lead  (Pb304),  or  potassium  per- 
manganate (KMnO4). 

Sometimes  chlorine  is  prepared  in  the  laboratory  by  heat- 
ing a  mixture  of  manganese  dioxide,  sodium  chloride,  and 
sulphuric  acid.  This  method  is  substantially  the  same  as 
the  other,  since  a  mixture  of  sulphuric  acid  and  sodium 
chloride  yields  hydrochloric  acid.  The  equation  for  this 
method  of  preparing  chlorine  is  — 

2  11,804  +  2  NaCl  +  MnO2  =  2  Cl  +  Na2SO4+MnSO4  +2  H2O 

Sulphuric        Sodium   Manganese   Chlorine      Sodium    Manganese     Water 
Acid  Chloride      Dioxide  Sulphate     Sulphate 

Chlorine  is  manufactured  by  several  processes,  one  of  which 
(Deacon  process)  involves  the  same  chemical  change  (oxidation)  as 
the  laboratory  method.  In  the  Deacon  process  hydrochloric  acid  is 
oxidized  by  oxygen  obtained  from  the  atmosphere.  A  mixture  of 
hydrochloric  acid  gas  and  air  is  heated  to  375°  C.  and  passed  through 
iron  tubes  containing  balls  of  clay  or  pieces  of  brick  previously  satu- 
rated with  copper  chloride  (CuCl2).  The  essential  chemical  change 
is  the  oxidation  of  the  hydrogen  of  the  hydrochloric  acid,  and  it  may 
be  represented  by  the  equation  — 

2HC1     +     O     =    2C1    +     H2O 

Hydrochloric      Oxygen       Chlorine          Water 
Acid 

The  copper  chloride  acts  as  a  catalytic  agent  in  this  process,  and  may 
be  replaced  by  other  chlorides. 

In  the  Weldon  process,  an  impure  native  manganese  dioxide,  known 
as  pyrolusite,  is  treated  with  hydrochloric  acid  in  large  earthenware 
retorts  or  stone  tanks  heated  by  hot  water  or  steam.  When  no  more 
chlorine  is  liberated,  the  residue  is  mainly  manganese  dichloride. 


198  INORGANIC  CHEMISTRY 

This  "still-liquor"  was  formerly  thrown  away,  but  by  the  Weldon 
process  it  is  changed  into  manganese  compounds,  which  are  used  to 
prepare  more  chlorine.  (See  Manganese  Dioxide.) 

Chlorine  is  also  prepared  on  a  large  scale  by  an  electrolytic  process. 
A  solution  of  sodium  chloride  is  electrolyzed  in  properly  constructed 
electrolytic  cells,'  and  the  chlorine  which  is  liberated  at  the  anode  is 
conducted  off  through  pipes.  Sodium  hydroxide  is  produced  at  the 
same  time,  and  the  process  will  be  described  under  this  compound. 

Properties.  —  Chlorine  is  a  greenish  yellow  gas.  Its  color 
suggested  the  name  chlorine  (from  the  Greek  word  chloros, 
meaning  greenish  yellow),  which  was  given  to  it  about  1810 
by  Davy,  who  spent  several  years  in  studying  this  gas  and 
its  compounds.  It  has  a  disagreeable,  suffocating  odor, 
which  is  very  penetrating.  If  breathed,  it  irritates  the 
sensitive  lining  of  the  nose  and  throat,  and  a  large  quantity 
would  doubtless  cause  death.  It  is  heavier  than  the  other 
elementary  gases,  and  is  about  2.5  times  heavier  than  air. 
Hence  it  is  easily  collected  by  conducting  it  to  the  bottom 
of  a  bottle  and  allowing  it  to  displace  the  air ;  the  term 
downward  displacement  is  sometimes  applied  to  this  method 
of  collecting  a  gas.  A  liter  of  dry  chlorine  at  0°  C.  and  760 
mm.  weighs  3.22  gm. 

Chlorine  is  moderately  soluble  in  water,  about  three  liters 
of  the  gas  dissolving  in  one  liter  of  water  under  ordinary 
conditions.  The  solution  is  yellowish,  smells  strongly  of 
chlorine,  and  is  frequently  used  in  the  laboratory  as  a 
substitute  for  the  gas.  Chlorine  water,  as  the  solution  is 
called,  is  unstable  even  under  ordinary  conditions,  and 
must  be  kept  in  the  dark.  If  the  solution  is  placed  in  the 
sunlight,  oxygen  is  soon  liberated  and  hydrochloric  acid  is 
formed.  Intermediate  changes  doubtless  occur;  but  the 
simplest  equation  for  the  essential  change  is  — • 

H2O     +     2C1    =     2HC1     +     O 

Water  Chlorine       Hydrochloric       Oxygen 

Acid 


CHLORINE  AND   HYDROCHLORIC  ACID          199 

Chlorine  is  much  less  soluble  in  a  solution  of  sodium  chloride, 
over  which  it  is  sometimes  collected.  It  attacks  mercury, 
and  cannot  be  collected  over  this  liquid. 

Chlorine  hydrate  is  formed  by  cooling  concentrated  chlorine  water, 
or  by  passing  chlorine  into  ice  water.  It  is  a  yellowish,  crystalline 
solid,  and  in  the  air  it  decomposes  quickly  into  chlorine  and  chlorine 
water.  Its  composition  corresponds  to  the  formula  Cl2 . 8  HaO. 

Liquid  Chlorine  was  first  prepared  by  Faraday  in  1823.  A  little 
chlorine  hydrate  was  inclosed  in  one  arm  of  a  bent  tube,  which  was 
then  sealed.  By  gently  heating  the  tube,  the  chlorine  hydrate  was 
decomposed  into  chlorine  and  water,  but  the  chlorine,  being  unable 
to  escape,  was  condensed  to  a  liquid  by  its  own  pressure  inside  the 
tube.  The  liquefaction  is  more  easily  accomplished,  if  one  end  is 
kept  cold  during  the  experiment.  This  method  has  been  replaced 
by  a  simpler  one;  viz.  subjecting  the  gas  to  a  high  pressure  and 
moderately  low  temperature.  The  critical  temperature  is  +141°  C. 
and  the  critical  pressure  is  84  atmospheres.  It  is  an  oily,  yellow 
liquid  and  is  now  a  common  commercial  substance.  Solid  chlorine 
is  a  yellow  crystalline  mass. 

Chlorine  is  an  active  chemical  element.  It  unites  directly 
with  most  of  the  elements,  the  only  conspicuous  exceptions 
being  oxygen,  carbon,  and  nitrogen.  Chlorine  does  not 
burn  in  the  air,  because  it  does  not  form  an  oxide  directly. 
Many  elements  unite  vigorously  with  chlorine.  Thus,  the 
metals  antimony  and  arsenic,  when  sprinkled  into  chlorine, 
suddenly  burst  into  flame,  while  the  non-metal  phosphorus 
melts  at  first  and  finally  burns  with  a  feeble  flame.  If 
sodium,  iron  powder,  copper  wire,  or  other  metals  are  heated 
and  then  put  into  chlorine,  they  burn;  the  sodium  and  iron 
produce  a  dazzling  light,  and  the  copper  glows  and  emits 
dense  fumes  of  whitish  smoke.  The  compound  formed 
in  each  case  is  a  chloride,  i.e.  a  compound  of  chlorine  and  one 
other  element.  (See  Chlorides,  below.)  Chlorine  combines 
readily  with  hydrogen.  Hence,  a  jet  of  burning  hydrogen 
when  lowered  into  chlorine  continues  to  burn,  forming 


200  INORGANIC  CHEMISTRY 

hydrochloric  acid  gas,  which  appears  as  a  white  cloud.  The 
simplest  equation  for  this  change  is  — 

H •     +     Cl     =     HC1 

Hydrogen    Chlorine    Hydrochloric 
Acid 

A  mixture  of  hydrogen  and  chlorine  explodes  violently  when 
exposed  to  the  sunlight.  Many  compounds  of  hydrogen 
are  decomposed  by  chlorine.  Thus,  compounds  of  hydrogen 
and  carbon,  such  as  those  found  in  illuminating  gas,  paraffin 
wax,  and  wood,  burn  in  chlorine  with  a  smoky  flame  ;  since 
chlorine  does  not  combine  directly  with  carbon,  the  flame 
contains  multitudes  of  very  fine  particles  of  solid  carbon. 
A  piece  of  glowing  charcoal  is  extinguished  by  chlorine. 
If  cotton  is  saturated  with  warm  turpentine  (Ci0H16)  and 
put  into  a  bottle  of  chlorine,  a  flame  accompanied  by  a  dense 
cloud  of  black  smoke  bursts  from  the  bottle ;  the  chlorine 
combines  with  the  hydrogen  to  form  hydrochloric  acid, 
while  the  carbon  is  left  free. 

The  power  to  bleach  is  the  most  striking  and  useful  prop- 
erty of  chlorine.  This  property  depends  upon  the  fact 
that  chlorine  and  water  interact  and  ultimately  liberate 
free  oxygen ;  the  latter  then  decomposes  the  complex  color- 
ing matter  into  colorless  substances.  If  an  envelope  on 
which  the  postmark,  or  a  pencil  mark,  is  still  visible  is  placed 
in  moist  chlorine,  these  marks  will  not  be  bleached  because 
they  are  largely  carbon;  the  writing  ink  will  disappear. 
Litmus  paper  and  many  kinds  of  colored  cloth  are  bleached 
by  moist  chlorine.  The  bleaching  action  of  chlorine  may  be 
explained  as  follows  :  Chlorine  and  water  form  hypochlorous 
acid  (HC1O),  and  then  this  very  unstable  acid  decomposes 
according  to  the  equation  — 

HC10       =     O     +     HC1 

Hypochlorous       Oxygen      Hydrochloric 
Acid  Acid 


CHLORINE  AND   HYDROCHLORIC  ACID          201 

Uses.  —  Chlorine  gas  is  used  extensively  to  manufacture 
bleaching  powder.  Liquid  chlorine  finds  application  in 
the  manufacture  of  chlorides  and  in  the  extraction  of  gold 
from  certain  ores.  Considerable  chlorine  gas  is  also  used 
to  bleach  wood  pulp;  both  gaseous  and  liquid  chlorine  are 
used  to  remove  bromine  from  its  compounds  on  a  large 
scale.  Bleaching  powder  is  used  as  a  germicide. 

Bleaching  Powder  is  the  main  source  of  the  chlorine  used 
in  the  bleaching  industries.  It  is  sometimes  called  "  bleach" 
or  "  chloride  of  lime."  It  is  a  yellowish  white  substance 
having  a  peculiar  odor,  which  resembles  that  of  chlorine. 
When  dry,  it  is  a  powder,  but  on  exposure  to  the  air,  it 
absorbs  water  and  carbon  dioxide,  becomes  lumpy  and 
pasty,  and  loses  some  of  its  chlorine,  owing  to  the  formation 
and  liberation  of  hypochlorous  acid  (HC1O).  Acids  like 
sulphuric  and  hydrochloric  acid  liberate  from  bleaching 
powder  its  "  available  chlorine/'  which  varies  from  30  to 
38  per  cent  in  good  qualities.  The  equations  for  the  inter- 
action of  acids  and  bleaching  powder  are  usually  written 
thus : — 


CaOCl2     +     H2SO4     =    2C1    +     CaS04     -f     H2O 

Bleaching  Sulphuric  Calcium 

Powder  Acid  Sulphate 

CaOCl2     +     2HC1     =    2C1    +     CaCl2       +     H2O 

Hydrochloric  Calcium 

Acid  Chloride 

The  composition  of  bleaching  powder  has  been  much  dis- 
cussed. The  most  reliable  authority  gives  it  the  formula 
CaOCl2.  When  dissolved  in  water,  bleaching  powder  forms 
calcium  hypochlorite  (CaO2Cl2)  and  calcium  chloride  (CaCl2). 

Bleaching  powder  is  manufactured  by  the  interaction  of  chlorine 
gas  and  lime.  Lime  (calcium  oxide,  CaO)  is  carefully  slaked  with 
water  to  form  calcium  hydroxide  (Ca(OH)2).  This  powder  is  sifted 


202  INORGANIC  CHEMISTRY 

into  a  large  absorption  chamber  made  of  iron,  lead,  or  tarred  brick  until 
the  floor  is  covered  with  a  layer  three  or  four  inches  deep.  The 
chlorine  enters  at  the  top  and  settles  slowly  to  the  floor.  The  simplest 
equation  for  the  formation  of  bleaching  powder  may  be  written  — 

Ca(OH)2     +    2C1    =     CaOCl2     +     H2O 

Calcium  Chlorine          Bleaching  Water 

Hydroxide  Powder 

Bleaching.  —  Immense  quantities  of  bleaching  powder  are 
used  to  whiten  cotton  and  linen  goods  and  paper  pulp.  The 
pieces  of  cotton  cloth  as  they  come  from  the  mill  are 
sewed  end  to  end  in  strips,  which  are  stamped  at  the  extreme 
ends  with  some  indelible  mark  to  distinguish  each  owner's 
cloth.  These  strips,  which  are  often  several  miles  long, 
are  drawn  by  machinery  into  and  out  of  numerous  vats  of 
liquors  and  water,  between  rollers,  and  through  machines, 
until  they  are  snow-white  and  ready  to  be  finished  (i.e. 
starched  and  ironed)  or  dyed.  The  whole  operation  re- 
quires three  or  four  days. 

The  preliminary  treatment  consists  in  singeing  off  the  downy  pile 
and  loose  threads  by  drawing  the  cloth  over  hot  copper  plates  or 
through  a  series  of  gas  flames.  The  object  of  the  remaining  opera- 
tions is  threefold:  (1)  to  wash  out  mechanical  impurities,  the  fatty 
and  resinous  matter,  and  the  excess  of  the  different  chemicals ;  (2)  to 
remove  matter  insoluble  in  water;  and  (3)  to  oxidize  the  coloring 
matter  by  chlorine.  The  details  of  the  process  differ  with  the  texture 
of  the  cloth  and  with  its  ultimate  use.  The  threefold  object  above 
mentioned  involves  successively  "liming,"  "souring,"  "chemicking," 
and  "souring,"  interspersed  with  frequent  washing.  The  "liming" 
consists  in  boiling  the  cloth  in  a  large  kier,  or  vat,  with  lime,  the 
"souring"  in  wetting  it  with  dilute  sulphuric  or  hydrochloric  acid, 
and  the  "chemicking"  in  impregnating  it  with  a  weak  solution  of 
bleaching  powder.  Often  the  cloth  is  boiled  at  a  certain  stage  with 
resin  and  sodium  carbonate.  The  "liming"  removes  the  resinous 
and  the  fatty  matter,  the  first  "souring"  neutralizes  traces  of  lime, 
and  the  second,  which  follows  the  "chemicking,"  liberates  the  chlorine 
in  the  fiber  of  the  cloth.  Frequent  washing  is  absolutely  necessary 
to  remove  the  impure  products  of  the  chemical  changes  as  well  as 


CHLORINE  AND   HYDROCHLORIC  ACID          203 

the  excess  of  lime  and  other  alkali,  acid,  and  chlorine.  Should  these 
be  left,  the  cloth  would  be  unevenly  bleached,  and  its  fiber  would  be 
weak.  The  cloth  is  finally  treated  with  an  antichlor,  such  as  sodium 
hyposulphite,  which  removes  the  last  traces  of  chlorine. 

Chlorides  are  formed  when  chlorine  combines  with  other 
elements,  and  they  are  in  general  stable  compounds.  The 
simplest  equations  illustrating  the  combination  of  chlorine 
with  certain  elements  are :  — • 


Na       + 

Cl 

=     NaCl 

Sodium 

Chlorine 

Sodium 
Chloride 

Sb      + 

3C1 

=     SbCl8 

Antimony 

Antimony 
Trichloride 

Cu      -f 

2C1 

=     CuCl, 

Copper 

Copper 
Chloride 

P         + 

3C1 

=     PC18 

Phosphorus 

Phosphorus 
Trichloride 

H      + 

Cl 

=       HC1 

Hydrogen 

Hydrochloric 
Acid 

Chlorides  are  an  important  class  of  compounds,  and  they  will 
be  considered  under  the  elements  with  which  chlorine  com- 
bines. (See  also  Chlorides,  below.) 

HYDROCHLORIC  ACID 

Hydrochloric  acid  is  the  common  name  of  a  water  solution 
of  a  very  important  compound  of  hydrogen  and  chlorine, 
viz.  hydrogen  chloride,  HC1.  Hydrogen  chloride  is  a  gas, 
which  is  very  soluble  in  water.  This  solution  is  known  in 
commerce  as  muriatic  acid  (from  the  Latin  word  muria, 
meaning  brine),  but  it  is  more  properly  called  hydrochloric 
acid.  Hydrogen  chloride  is  often  called  hydrochloric  acid 
gas. 


204  INORGANIC  CHEMISTRY 

Occurrence.  —  Hydrogen  chloride  occurs  free  in  volcanic 
gases.  The  solution  is  one  constituent  of  the  gastric  juice. 

Preparation.  —  The  gas  is  prepared  by  heating  concen- 
trated sulphuric  acid  and  sodium  chloride  in  an  apparatus 
like  that  used  for  chlorine  (Fig.  23).  If  the  mixture  is  gently 
heated,  the  chemical  change  is  represented  thus :  - 

NaCl     +     H2SO4     =     HC1     +     HNaSO4 

Sodium  Sulphuric        Hydrochloric        Acid  Sodium 

Chloride  Acid  Acid  Sulphate 

But  at  a  high  temperature  the  equation  for  the  reaction  is  — 
2  NaCl     +     H2SO4     =     2  HC1     +     Na2SO4 

Sodium 
Sulphate 

The  solution  is  prepared  by  passing  the  gas  into  water. 

If  sulphuric  acid  is  added  to  a  solution  of  sodium  chloride 
instead  of  the  solid,  little  or  no  hydrogen  chloride  is  liberated 
(unless  heat  is  applied).  This  is  due  to  the  fact  that  the 
hydrogen  chloride  being  very  soluble  in  water  remains  in  the 
sphere  of  action  and  tends  to  cause  the  equilibrium  — 

NaCl  +  H2S04  ;£  HC1  +  HNaSO4 

It  has  already  been  pointed  out  that  one  way  to  displace  equi- 
librium is  to  remove  a  product  of  the  reaction.  Thus,  in  this 
case  the  volatile  hydrogen  chloride  escapes  from  the  mixture 
of  sulphuric  acid  and  (solid)  sodium  chloride  and  allows  the 
forward  reaction  to  proceed. 

Commercial  Hydrochloric  Acid  is  manufactured  by  heat- 
ing a  mixture  of  salt  and  sulphuric  acid  to  a  moderate 
temperature  in  a  hemispherical  cast  iron  retort,  and  conduct- 
ing the  gas  through  an  earthenware  pipe  into  an  absorbing 
tower ;  the  fused  mass  of  acid  sodium  sulphate  and  salt  is 
then  subjected  to  a  higher  temperature,  and  the  remainder 
of  the  gas  is  led  into  the  absorbing  tower.  These  towers  are 


CHLORINE  AND  HYDROCHLORIC  ACID          205 

high  and  filled  with  coke  or  pieces  of  brick  over  which  water 
trickles  ;  as  the  hydrochloric  acid  gas  passes  up  the  tower, 
it  is  absorbed  by  the  descending  water,  and  concentrated 
acid  flows  from  the  bottom  of  the  tower.  The  gas  is  usually 
cooled  before  it  enters  the  towers.  Sometimes  the  gas  is 
conducted  through  huge  earthenware  jars  before  entering 
the  towers.  In  these  jars  the  gas  and  water  are  caused  to 
flow  constantly  in  opposite  directions,  thus  insuring  com- 
plete absorption.  Hydrochloric  acid  can  be  manufactured 
synthetically,  i.e.  by  burning  hydrogen  gas  in  chlorine 
gas.  The  equation  for  this  reaction  is  - 
H  +  Cl  =  HC1 

Hydrogen     Chlorine     Hydrochloric 
Acid 

Properties.  —  Hydrochloric  acid  gas  is  colorless.  When 
it  escapes  into  moist  air,  it  forms  fumes  which  are  really 
minute  drops  of  a  solution  of  the  gas  in  the  moisture  of  the 
air.  It  has  a  choking,  sharp  taste,  and  irritates  the  lining 
of  the  nose  and  throat.  The  gas  does  not  burn  nor  support 
combustion.  It  is  about  1.25  times  heavier  than  air,  and 
may  therefore  be  collected  by  displacement  of  air,  like 
chlorine.  One  liter  at  0°  C.  and  760  mm.  weighs  1.641  gm. 
The  critical  temperature  is  4-52°  C.,  and  the  gas  becomes 
a  colorless  liquid  when  subjected  to  a  high  pressure  and 
moderately  low  temperature.  The  extreme  solubility  of 
hydrochloric  acid  gas  in  water  is  one  of  its  most  striking 
properties.  One  liter  of  water  will  dissolve  about  550  1. 
of  gas,  if  both  are  at  0°  C.  and  760  mm.  At  the  ordinary 
temperature  about  500  1.  of  gas  dissolve  in  1  1.  of  water, 
and  as  the  temperature  rises,  the  solubility  decreases.  The 
solution  is  the  familiar  hydrochloric  acid.  The  gas  readily 
escapes,  hence  the  acid  forms  fumes  when  exposed  to  the 
air.  Pure  hydrochloric  acid  is  a  colorless  liquid,  but  the 
commercial  acid  has  a  yellow  color,  usually  due  to  iron 


206  INORGANIC  CHEMISTRY 

compounds  or  to  dissolved  chlorine.  Like  the  common 
acids,  its  solution  reddens  blue  litmus  and  yields  hydrogen 
by  interaction  with  certain  metals.  In  terms  of  the  theory 
of  electrolytic  dissociation,  hydrochloric  acid  is  a  strong 
acid;  i.e.  it  dissociates  to  a  considerable  extent  into  ions, 
one  kind  being  H+,  the  other  being  Cl~. 

The  most  concentrated  acid  contains  about  40  per  cent 
(by  weight)  of  the  compound  (HC1),  and  its  specific  gravity 
is  1.2.  When  the  concentrated  acid  is  heated,  the  gas  is 
evolved  until  the  solution  contains  about  20  per  cent  of  the 
acid,  and  then  the  liquid  boils  at  110°  C.  without  further 
change  in  concentration.  The  dilute  acid  loses  water  until 
the  same  conditions  prevail. 

Uses  of  Hydrochloric  Acid.  —  Vast  quantities  are  used 
to  prepare  the  chlorine  consumed  in  the  manufacture  of 
bleaching  powder.  Various  chlorides  are  prepared  from 
it,  and  it  is  one  of  the  common  acids  used  in  the  chemical 
laboratory  and  in  many  industries. 

Chlorides  are  formed  by  direct  addition  of  chlorine  to 
metals,  as  we  have  seen.  They  are  also  formed  by  the 
substitution  of  a  metal  for  the  hydrogen  in  hydrochloric 
acid.  Chlorides,  therefore,  are  salts  of  hydrochloric  acid. 
They  can  be  prepared  in  several  ways,  e.g.  (1)  by  the  inter- 
action of  hydrochloric  acid  and  metals,  metallic  oxides, 
or  hydroxides,  and  (2)  by  the  interaction  of  certain  salts 
and  hydrochloric  acid  or  chlorides.  The  following  equa- 
tions illustrate  several  of  these  methods  :  — 

Zn     -f-     2HC1     =     ZnCls     +     2  H 

Zinc  Zinc 

Chloride 

ZnO     H-     2HC1     =     ZnCl2     +     H2O 

Zinc  Zinc 

Oxide  Chloride 


CHLORINE  AND   HYDROCHLORIC  ACID          207 
Zn(OH)2     +     2HC1     =     ZnCl2     +     2  H2O 

Zinc  Zinc 

Hydroxide  Chloride 

AgNOa      +      HC1      =     AgCl      +     HN03 

Silver  Silver  Nitric 

Nitrate  Chloride  Acid 

One  molecule  of  a  chloride  may  contain  several  atoms  of 
chlorine.  Often  the  name  of  the  compound  indicates  this 
fact;  e.g.  manganese  dichloride  (MnCl2),  antimony  tri- 
chloride (SbCl3),  phosphorus  trichloride  and  pentachloride 
(PC13  and  PC15).  If  a  metal  forms  two  chlorides,  the  two 
are  distinguished  by  modifying  the  name  of  the  metal; 
the  one  containing  the  smaller  proportions  of  chlorine  ends 
in  -ous,  the  one  containing  the  larger  ends  in  -ic.  Thus, 
mercurous  chloride  is  HgCl,  but  HgCl2  is  mercuric  chloride. 
Similarly,  we  have  ferrous  chloride,  FeCl2,  and  ferric  chloride, 
FeCl8. 

The  Test  for  Hydrochloric  Acid  and  Chlorides.  —  Most 
chlorides  are  soluble  in  water,  those  of  lead,  silver,  and 
mercury  (-ous)  being  the  only  conspicuous  exceptions. 
If  silver  nitrate  is  added  to  hydrochloric  acid  or  to  the 
solution  of  a  chloride,  a  white,  curdy  precipitate  of  silver 
chloride  is  formed,  which  (a)  is  insoluble  in  nitric  acid, 
(6)  soluble  in  warm  ammonium  hydroxide,  and  (c)  turns 
purple  in  the  sunlight.  The  invariable  formation  of  silver 
chloride  is  the  test  for  hydrochloric  acid  and  soluble  chlorides. 
This  chemical  change  is  a  typical  illustration  of  double  de- 
composition. The  equation  for  the  chemical  change  is  — • 

HC1     +     AgNO3-    =     AgCl     +     HN08 

Hydrochloric  Silver  Silver  Nitric 

Acid  Nitrate  Chloride  Acid 

By  an  inspection  of  this  equation  we  see  that  both  the  hydro- 
chloric acid  and  silver  nitrate  decompose  and  the  chemical 
fragments,  so  to  speak,  recombine  in  a  different  way  to 


208  INORGANIC   CHEMISTRY 

form  the  two  products.  The  same  test  is  applicable  to 
hydrochloric  acid  and  a  chloride,  because  both  yield  chlorine 
ions  (Cl~).  The  test  is  for  ionic  chlorine,  and  the  general 
ionic  equation  may  be  written  thus  :  — 


Miscellaneous.  —  Besides  hydrochloric  acid,  there  are 
four  other  acids  of  chlorine  ;  they  contain  oxygen  as  well 
as  hydrogen.  These  acids  are  hypochlorous  acid  (HC1O), 
chlorous  acid  (HC1O2),  chloric  acid  (HC1O3),  and  perchloric 
acid  (HC1O4).  They  are  prepared  with  difficulty,  and  de- 
compose readily.  Thus,  hypochlorous  acid  is  prepared 
by  the  interaction  of  dilute  hydrochloric  acid  and  sodium 
hypochlorite  (NaCIO),  but  the  procedure  must  be  careful, 
because  the  resulting  hypochlorous  acid  itself  is  decom- 
posed by  hydrochloric  acid.  Hypochlorous  acid  also  de- 
composes when  its  solution  is  warmed  or  exposed  to  the 
sunlight,  oxygen  gas  and  hydrochloric  acid  being  the  prod- 
ucts. This  acid  (HC1O)  is  formed  to  a  slight  extent  when 
chlorine  dissolves  in  water,  thus  :  - 

2C1     +     H2O     =     HC1O     +     HC1 

Chlorine  Water        Hypochlorous     Hydrochloric 

Acid  Acid 

But  the  hypochlorous  acid  decomposes  readily  and  under 
many  conditions,  especially  when  exposed  to  the  sunlight; 
the  equation  for  the  reaction  in  the  last-named  case  may  be 
written  :  — 

HC1O     =     O     +     HC1 

Hypochlorous      Oxygen     Hydrochloric 
Acid  Acid 

It  is  this  free  oxygen  which  is  the  effective  agent  in  bleach- 

ing,  though   the  unstable   hypochlorous   acid    is    the   way 

station,  so  to  speak,  between  the  chlorine  and  the  oxygen. 

The  corresponding  salts  of  two  of  the  oxy-chlorine  acids 


CHLORINE  AND  HYDROCHLORIC  ACID         209 

are  important,  viz.  the  hypochlorites  and  chlorates.  Solu- 
tions of  potassium  hypochlorite  (Javelle's  water)  and  sodium 
hypochlorite  (Labarraque's  solution)  find  application  in  the 
removal  of  stains  from  cotton  and  linen  goods ;  "  bleach 
liquors "  consisting  largely  of  hypochlorites  are  use  as 
bleaching  agents  in  some  industries.  Chlorates  of  potassium 
and  sodium  are  used  as  a  source  of  oxygen. 

Chlorine  forms  two  oxides  —  chlorine  monoxide  (C^O) 
and  chlorine  dioxide  (C102).  Each  is  an  unstable  yellowish 
brown  gas. 

PROBLEMS  AND  EXERCISES 

1.  Calculate  the  volume  of  hydrochloric  acid  solution  (having 
a  specific  gravity  of  1.10  and  containing  20.12  per  cent  of  HC1) 
required  to  make  500  gm.  of  barium  chloride  from  barium  carbonate. 

2.  What  weight  of  chlorine  can  be  prepared  from  78  gm.  of  rock 
salt,  containing  99  per  cent  of  NaCl?     Howmuch  manganese  diox- 
ide and  sulphuric  acid  would  be  required?     What  volume  would 
the  chlorine  occupy  under  standard  conditions  ? 

3.  (a)  Magnesium  chloride  heated  in  steam  forms  hydrogen 
chloride  and  magnesium  oxide.     Write  the  equation.     (6)  What 
volume  of  hydrogen  chloride  (at  0°  C.  and  760  mm.)  can  be  formed 
from  50  gm.  of  magnesium  chloride? 

4.  (a)  Calcium  oxide  and  hydrochloric  acid  form  calcium  chlo- 
ride and  water.     Write  the  equation.     (6)  What  volume  of  hydro- 
gen chloride  (at  0°  C.  and  760  mm.)  is  needed  to  produce  50  gm. 
of  calcium  chloride? 

5.  Write  equations  for  the  following  reactions  by  applying  the 
method  outlined  in  Chapter  VII :  (a)  Phosphorus  and  chlorine  form 
phosphorus  trichloride.     (6)  Phosphorus   trichloride  and  chlorine 
form  phosphorus  pentachloride.     (c)  Aluminium  and  hydrochloric 
acid  form  hydrogen  and  aluminium  chloride. 

6.  Write  ionic  equations  for  the  interaction  of  (a)  hydrochloric 
acid  and  silver  sulphate,  and  (6)  calcium  chloride  and  sodium  sul- 
phate. 

7.  Interpret  the  changes  in  6  from  the  standpoint  of  solubility 
product. 

8.  Interpret   the  preparation   of  hydrogen   chloride  from   the 
standpoint  of  displacement  of  equilibrium. 


CHAPTER  XIII 
Compounds  of  Nitrogen  —  Gay-Lussac's  Law 

THE  most  important  compounds  of  nitrogen  ^ire  ammonia 
(NH3),  nitric  acid  (HNO3),  and  compounds  related  to  them. 
Many  animal  and  vegetable  substances  essential  to  life  are 
compounds  of  nitrogen. 

AMMONIA 

The  term  ammonia  includes  both  the  gas  and  its  solution 
in  water,  though  the  latter  is  more  accurately  called  am- 
monium hydroxide. 

Formation  of  Ammonia.  —  When  vegetable  and  animal 
matter  containing  nitrogen  decomposes  or  decays,  the 
nitrogen  and  hydrogen  are  liberated  in  combination  as 
ammonia.  The  odor  of  ammonia  is  often  noticed  near 
stables.  If  animal  matter  containing  nitrogen  is  heated, 
ammonia  is  given  off.  The  old  custom  of  preparing  ammonia 
by  heating  horns  and  hoofs  in  a  closed  vessel,  i.e.  by  dry 
distillation,  gave  rise  to  the  term  "  spirits  of  hartshorn." 
Soft  coal  contains  compounds  of  nitrogen  and  of  hydrogen, 
and  when  the  coal  is  heated  to  make  illuminating  gas,  one 
of  the  products  is  ammonia. 

Preparation.  —  Ammonia  gas  is  prepared  in  the  laboratory 
by  heating  ammonium  chloride  with  a  base,  the  mild  base 
calcium  hydroxide  being  usually  used.  The  equation  for 
the  reaction  is  — 

2NH4C1     +     Ca(OH)2     =     2NH4OH     +      CaCl2 

Ammonium  .Calcium  Ammonium  Calcium 

Chloride  Hydroxide  Hydroxide  Chloride 

210 


COMPOUNDS   OF   NITROGEN 


211 


But  the  ammonium  hydroxide  is  unstable  and  decomposes 
into  ammonia  gas  and  water,  thus  :  — 


NH4OH 

Ammonium 
Hydroxide 


NH8 

Ammonia 
Gas 


H20 

Water 


The  ammonia  gas  is  very  volatile  and  is  usually  collected 

by  allowing  it  to  flow  into  an  inverted  bottle  and  displace 

the  air  (Fig.   24) ;   the  term  upward 

displacement  is  sometimes  applied  to 

this  method  of  collecting  a  gas.     The 

solution    is   prepared   by    conducting 

the  gas  into  water. 

The  main  source  of  the  ammonia  of  com- 
merce is  the  ammoniacal  liquor  or  gas  liquor 
of  the  illuminating  gas  works.  The  gases 
which  come  from  the  retorts  in  which  the 
coal  is  heated  are  passed  into  water,  which 
absorbs  the  ammonia  and  certain  other 
gases.  This  impure  gas  liquor  is  treated 
with  an  alkali  to  liberate  the  ammonia,  which 
is  absorbed  in  tanks  containing  hydrochloric 
acid  or  sulphuric  acid.  This  solution  upon  FlG-  24.  —  Apparatus  for 
the  addition  of  an  alkali  (such  as  calcium  collecting  ammonia  gas. 
i  ,  .  ,  x  .  . ,  .  ,  .  ,  .  The  gas  from  the  gen- 

hydroxide)  gives  up  its  ammonia,  which  is       erator  flowg  througlf  the 

dissolved  in  distilled  water,  forming  thereby  delivery  tube  B  to  the 
the  ammonium  hydroxide  or  aqua  ammoniae  top  of  the  bottle  D  and 
of  commerce.  displaces  the  air,  which 

escapes  through  the  hole 

Properties  of  Ammonia.  —  Ammonia      in  the  block  Cm 
gas  is   colorless.     It   has   an  exceed- 
ingly pungent   odor,   and   if  inhaled   suddenly  or  in  large 
quantities  it  brings  tears  to  the  eyes  and  may  cause  suf- 
focation.    It  is  a  light,  volatile  gas,  being  only  .59  times 
as  heavy  as  air.     A  liter  of  the  gas  at  0°  C.  and  760  mm. 
weighs  about  .77  gm.     It  will  not  burn  in  the  air  under 
ordinary   conditions,   nor  will  it    support  the   combustion 


212  INORGANIC  CHEMISTRY 

of  a  blazing  stick  ;  but  if  the  air  is  heated,  or  if  considerable 
oxygen  is  mixed  with  the  air,  then  a  jet  of  ammonia  gas 
may  be  made  to  burn  (in  a  suitable  apparatus)  with  a  pale 
yellowish  flame. 

Ammonia  gas  is  easily  liquefied,  since  its  critical  tem- 
perature is  about  +130°  C.;  at  0°  C.  a  pressure  of  about 
4.2  atmospheres  causes  liquefaction.  At  about  —76°  C. 
a  white  solid  is  produced.  Liquefied  ammonia  is  often 
called  anhydrous  ammonia,  because  it  contains  no  water. 
It  boils  at  about  —  33.5°  C.  Hence,  if  it  is  exposed  to 
the  air  or  warmed  in  any  way,  it  changes  back  to  a  gas, 
and  in  doing  so  absorbs  considerable  heat.  This  fact  has 
led  to  the  extensive  use  of  liquid  ammonia  in  the  manu- 
facture of  ice. 

When  electric  sparks  are  passed  through  ammonia  gas, 
it  decomposes  into  nitrogen  and  hydrogen.  But  if  a  mixture 
of  nitrogen  and  hydrogen  is  sparked,  these  gases  combine 
to  form  ammonia  gas.  In  neither  case,  however,  is  the  re- 
action complete.  The  final  mixture  always  contains  about 
98  per  cent  of  nitrogen  and  hydrogen,  and  2  per  cent  of 
ammonia  gas.  That  is,  the  reaction  is  reversible  and  pro- 
ceeds until  equilibrium  is  reached.  The  simplest  equation 
for  this  reversible  reaction  is  — 


If  water  or  acid  is  added  during  the  sparking  of  the  mixture 
of  nitrogen  and  hydrogen,  the  ammonia  is  dissolved.  Its 
removal  displaces  the  equilibrium,  and  the  reaction  pro- 
ceeds to  completion;  i.e.  all  the  nitrogen  and  hydrogen 
combine. 

Another  marked  property  of  ammonia  gas  is  its  solubility 
in  water.  A  liter  of  water  at  0°  C.  dissolves  1148  1.  of  gas 
(measured  at  0°  C.  and  760  mm.),  and  at  the  ordinary  tem- 
perature 1  1.  of  water  dissolves  about  700  1.  of  gas.  This 


COMPOUNDS   OF   NITROGEN  213 

solution  of  the  gas  is  usually  called  ammonia,  though  other 
names  are  often  applied  to  it;  i.e.  ammonium  hydroxide, 
aqua  ammonias,  or  ammonia  water.  It  gives  off  the  gas  freely, 
especially  when  heated,  as  may  easily  be  discovered  by  the 
odor  or  by  the  formation  of  the  dense  white  fumes  of  solid 
ammonium  chloride  (NH4C1)  when  the  solution  is  exposed 
to  hydrochloric  acid.  The  volatility  of  ammonia  was  early 
detected,  and  the  name  volatile  alkali  was  applied  to  it; 
the  discoverer,  Priestley,  called  the  gas  alkaline  air.  The 
gas  can  be  completely  removed  from  solution  by  boiling. 
The  commercial  solution  is  lighter  than  water  (its  specific 
gravity  being  about  .88)  and  contains  about  35  per  cent 
(by  weight)  of  the  gas.  Ammonium  hydroxide  has  an  alka- 
line reaction,  neutralizes  acids  and  forms  salts,  and  acts 
in  many  respects  like  sodium  hydroxide.  In  terms  of  the 
theory  of  electrolytic  dissociation,  ammonium  hydroxide 
is  a  weak  base ;  i.e.  it  dissociates  only  to  a  slight  degree  (1.4 
per  cent  in  N/10  solution  at  18°  C.)  into  ions  (NH4+  and  OH~). 

Ammonium  Compounds.  —  When  ammonia  gas  is  passed 
into  water,  some  of  the  gas  combines  with  the  water  and 
forms  a  solution  of  an  unstable  compound  called  ammonium 
hydroxide.  Its  formation  may  be  represented  thus  :  — 

NH3     +     H2O     =     NH4OH 

Ammonia  Water  Ammonium 

Hydroxide 

Ammonia  (NH3)  also  unites  directly  with  acids,  thereby 
forming  salts,  thus  :  — 

NH3     +     HC1     =      NH4C1 

Ammonia     Hydrochloric       Ammonium 
Gas  '    Acid  Chloride 

Analogous  salts  are  ammonium  sulphate  and  ammonium 
nitrate. 


214  INORGANIC   CHEMISTRY 

Ammonium  hydroxide  is  a  base.     It  neutralizes  acids  and 
forms  salts,  thus  :  — 

NH4OH     +     HC1     =     NH4C1     +     H2O 

Ammonium 
Chloride 

2NH4OH     +     HaSO,     =     (NH4)2SO4     +     2H2O 

Ammonium 
Sulphate 

These  salts,  ammonium  chloride  and  ammonium  sulphate, 
are  strictly  analogous  to  sodium  salts.     Thus,  we  have  - 

SODIUM  SALTS  AMMONIUM  SALTS 

NaCl  NH4C1 

NaNO3  NH4NO3 


etc.  etc. 

Hence,  it  is  believed  that  ammonium  compounds  contain 
a  group  of  atoms  which  acts  chemically  like  an  atom  of  a 
metal.  This  group  of  atoms  is  called  ammonium,  and  its 
formula  is  NH4.  Ammonium  has  never  been  separated 
from  its  compounds,  or  if  it  has,  it  is  so  unstable  that  it 
immediately  decomposes  into  ammonia  gas  and  hydrogen. 
So  also  ammonium  hydroxide  has  never  been  obtained  free, 
for  it  decomposes  readily  into  ammonia  gas  and  water.  How- 
ever the  properties  of  ammonium  hydroxide  leave  no  doubt 
but  that  it  is  a  compound  of  hydro  xyl  and  ammonium. 
Ammonium  is  called  a  radical,  because  it  is  the  root  of  a 
series  of  compounds. 

Uses  of  Ammonium  Hydroxide.  —  Ammonium  hydroxide 
is  widely  used  as  a  cleansing  agent  (especially  for  .the  removal 
of  grease),  as  a  restorative  in  case  of  fainting  or  of  inhaling 
irritating  gases,  in  dyeing,  and  calico  printing,  and  in  the 
manufacture  of  dyestuffs,  sodium  bicarbonate,  and  am- 
monium compounds. 


COMPOUNDS  OF  NITROGEN 


215 


The  Use  of  Ammonia  as  a  Refrigerant  and  in  making  Ice 

depends  on  the  fact  that  liquefied  ammonia  (not  ammonia 
water)  changes  rapidly  into  a  gas  and  thereby  absorbs  heat. 
Hence,  if  liquefied  ammonia  is  allowed  to  flow  through  a  pipe 
immersed  in  a  solution  of  sodium  chloride  or  calcium  chloride 
(technically  called  a  brine),  the  ammonia  evaporates  in  the 
pipe  and  cools  the  brine,  which  may  be  used  directly  as  a  re- 
frigerant or  for  making  ice.  In  some  cold  storage  plants, 
breweries,  packing  houses,  and  sugar  refineries  this  cold  brine 
is  circulated  through  a  system  of  pipes  placed  in  the  rooms 
where  a  low  temperature  is  desired. 

Cold  Water  trickling  over  the 
ammonia  pipes  to  condense 


the  compressed  gas 


Ammonia  pump 
Low  Pressure  3T        High  Pressure 


Expansion  valve 
Brine  pump 

Fig.  24  a.  —  Diagram  of  an  ice-making  plant. 

In  the  operation  of  an  ice-making  plant  (Fig.  24  o)  liquefied 
ammonia  is  forced  from  a  tank  into  a  series  of  pipes  which  are 
submerged  in  an  immense  vat  filled  with  brine.  Large  galvanized 
iron  cans  containing  pure  water  to  be  frozen  are  immersed  in  the 
brine,  which  is  kept  below  the  freezing  point  of  water  by  rapid 
evaporation  of  the  ammonia  in  the  pipes.  After  several  hours  the 
water  in  the  cans  is  frozen  into  ice.  As  fast  as  the  ammonia  gas 
forms  in  the  pipes,  it  is  removed  by  exhaust  pumps  into  another  tank, 
where  it  is  recondensed  to  liquefied  ammonia  and  conducted,  as 
needed,  into  the  first  tank  ready  for  renewed  use.  The  ammonia  is 
thus  used  over  and  over  without  appreciable  loss.  The  pure  water 
is  often  obtained  from  an  artesian  well.  Most  ocean  steamers  have 
an  ice  plant  and  cold  storage  room ;  in  cities  in  warm  climates 
manufactured  ice  is  a  common  commodity. 


216  INORGANIC   CIIKMISTKY 

Composition  of  Ammonia  Gas.        Many  facts  show  that  ammonia 
gas  is  composed  of    nitrogen  and    hydrogen  in  the  ratio  of  1   to  '•'>    \>y 
volume  and  14  to  li  by  weight.     The  volumetric  com- 
position is  shown  in  two  ways.    The  first,  depends  upon 
the   fact    thai    ammonia,  and   chlorine  interact    thus:  — 

2NH8     +     3C12    =     Nj,    +     OIK 'I 

Ammonia  CUmim-        Nitroj/rn       1 1  \  <ln>rlilojir 

(ias  'Acid 

A  tube    (Fig.  2f>)   filled   with    a    known    volume   of 
chlorine  is  provided  with  a  funnel   through  which  con 
centraled    ammonium    hydroxide    is   dropped    into    the 
chlorine,  until  the  reaction  oe&fieft.      After  the  excess  of 
ammonia  is  neiit  rali/.ed  with  sulphuric  acid,  the  volume 
of  nitrogen  left  is  found  to  be  one  third  of  the  original 
volume  of  chlorine  gas.      Now   hydrogen    and    chlorine 
combine  in  e<|iial  volumes;   hence,  the  volume  of  hydro- 
gen   withdrawn    from    the    added    ammonia  must    l>e 
Fia. 25. —  Ap-  equal    to   the  original  volume  of  chlorine.     But    this 
p:n-:ihiH    for   volume  is  three?  times  the  volume  of  nitrogen;  therefore 
determining    (|M.n.  js  t,,m.  jim(.s  a8  much  hydrogen  as  nitrogen  by 
'.".'  """I"'"   volume  in  ammonia  gas. 

Ml  ion  nl  ;mi  -  ...  tA.tr  i 

nioni      aH  gravimetric   composition  of    ammonia  g;is    is 

found  by  oxidizing  it,  and  weighing  the  water  and 
nitrogen,  which  are  the  only  products.  The  results  show  that  four- 
teen parts  of  nitrogen  combine  with  three  parts  of  hydrogen. 

NITRIC  ACID 

Nitric  Acid  is  one  of  the  most  useful  compounds  of  nitro- 
gen. It  was  known  to  the  alchemists,  who  used  it  to  pre- 
pare a  mixture  which  dissolves  gold.  (See  Aqua  Regin ,  below.) 
Nitric  acid  is  used  in  the  preparation  of  many  nitrogen  com- 
pounds. 

Formation  of  Nitric  Acid.  —  When  moist  animal  or  vege- 
table matter  containing  nitrogen  decays  in  the  presence 
of  an  alkali,  nitric  acid  is  formed  ;  but  it  is  neutralized  at 
once  by  the  alkali,  so  nitrates  — salts  of  nitric  acid  —  are  the 
final  products.  This  chemical  change  is  known  as  nitri- 


COMPOUNDS  OF  NITROGEN  217 

fication,  ami  it  in  caused,  or  largely  influen^  -d,  \>y  minute; 
Jiving  organisms  called  bacteria. 

Nitric  ;ifid  is  formed  when  electric  sparks  are  passed 
through  moist  air.  The  nitrogen  and  oxygen  <-<>ut\>\n<-  to 
-orn<-  extent  and  form  nitric  oxid'-  (NO;,  whirh  unites  with 
oxygen  find  form-  nitrogen  dioxide  (NO2).  This  gas  and 
water  form  nitric  acid  thus  :  — 

3  NO2  +  HjO  -  2  HNOa  +  NO 

Nitrogen        Water  Ni'-  Nit' 

Dioxide  Add  Oxide 

This  chemical  change  is  now  being  applied  oa  a  commercial 
scale.  Air  is  forced  through  a  tube  in  which  a  powerful 
electric  arc  is  spread  out  into  a  disk  by  an  electromagnet. 
The  nitrogen  dioxide  is  absorbed  in  water  or  in  a  mixture  of 
water  and  lime,  thereby  forming  nitric  acid  or  calcium  nitrate. 

Preparation.  —  Nitric  acid  is  prepared  in  the  laboratory 
by  heating  concentrated  sulphuric  acid  with  a  nitrate 
in  a  glass  retort  ;  the  nitric  acid  distills  into  a  receiver,  which 
is  kept  cool  by  water.  The  reaction  at  ordinary  temperatures 
is  represented  by  the  equation  — 

NaNO,  +  H2S04  ;£  HNO,  +  HNa8O4 

Hodhim  Sulphuric  Nitric  Ad/1  Sodium 

Nitrate  Acid  Acid 


The  reaction  is  reversible  and  equilibrium  is  noon  established. 
But  since  nitric  acid  boils  at  86°  C.,  it  is  removed  by  gentle 
heat  and  recovered  by  condensation,  It»  removal  displaces 
the  equilibrium  and  allows  the  forward  reaction  to  proceed 
to  completion.  At  a  high  temperature  sodium  sulphate 
(Na^O4)  is  formed,  but  since  part  of  the  nitric  add  is  de- 
composed, excessive  heat  is  avoided. 

Nitric  acid  is  manufactured  by  heating  sodium  nitrate 
and  sulphuric  add  in  a  large  cast-iron  retort  and  condensing 


218  INORGANIC  CHEMISTRY 

the  vapors  in  huge  glass  or  earthenware  bottles;  the  last 
bottle  is  connected  with  a  tower  filled  with  coke  over  which 
water  trickles  to  absorb  the  vapors  which  escape  from  the 
bottles.  The  acid  vapors  are  also  often  condensed  in  earthen- 
ware or  glass  tubes. 

Properties.  —  Pure  nitric  acid  is  a  colorless  liquid,  but  the 
commercial  acid  is  yellow  or  reddish,  owing  to  absorbed 
nitrogen  compounds,  chlorine,  or  iron  compounds.  It 
decomposes  slowly  in  the  sunlight  or  when  heated,  and 
a  brownish  gas  may  often  be  seen  in  bottles  of  nitric  acid. 
It  absorbs  water,  and  forms  irritating  fumes  when  exposed 
to  the  air.  The  specific  gravity  of  the  commercial  acid  is 
about  1.42,  and  it  contains  approximately  70  per  cent  of  the 
real  acid  (HNOs),  the  rest  being  water.  If  the  water  is 
removed  by  slowly  distilling  the  commercial  acid  with  con- 
centrated sulphuric  acid,  the  product  contains  from  94  to 
99  per  cent  of  the  real  acid,  and  its  specific  gravity  is  about 
1.51.  When  nitric  acid  is  boiled,  it  loses  either  acid  or  water 
until  the  liquid  contains  approximately  68  per  cent  of  nitric 
acid,  and  then  it  continues  to  boil  unchanged  in  concentration 
at  120°  C.  (See  Hydrochloric  Acid.) 

A  solution  of  nitric  acid  has  the  properties  of  acids  to 
a  marked  degree.  In  terms  of  the  theory  of  electrolytic 
dissociation  it  is  one  of  the  strongest  acids ;  i.e.  it  dissociates 
largely  into  ions,  one  kind  being  H+,  the  other  being  NO3~. 

Nitric  acid  is  very  corrosive.  It  turns  protein,  e.g.  the  skin, 
a  yellow  color  owing  to  the  formation  of  xanthoprotein ;  the 
concentrated  acid  causes  serious  burns  and  should  be  used 
with  extreme  care.  Nitric  acid  decomposes  readily,  espe- 
cially when  hot,  and  is  therefore  an  energetic  oxidizing  agent. 
Glowing  charcoal  continues  to  burn  in  the  acid,  while  straw, 
sawdust,  hair,  and  similar  substances  are  charred  and  even 
inflamed  by  it.  Iron  sulphide  (FeS)  heated  with  nitric  acid 


COMPOUNDS  OF  NITROGEN  219 

becomes  oxidized  to  iron  sulphate  (FeSO4).  It  interacts 
readily  and  often  violently  with  metals,  metallic  oxides, 
and  hydroxides,  forming  a  variety  of  products,  especially 
nitrates. 

Uses  of  Nitric  Acid.  —  Nitric  acid  is  one  of  the  common 
laboratory  acids.  Large  quantities  are  used  in  the  manu- 
facture of  nitrates,  dyestuffs,  sulphuric  acid,  nitroglycerin, 
and  guncotton,  and  in  the  refining  of  gold  and  silver. 

Nitrates.  —  Nitric  acid  is  monobasic  and  forms  a  series  of 
well-defined  salts  called  nitrates.  The  interaction  of  nitric 
acid  and  most  metals  is  exceedingly  vigorous,  and  for  this 
reason,  probably,  the  alchemists  called  the  acid  aqua  fortis 
—  strong  water.  The  reaction  varies  with  the  metal,  con- 
centration of  the  acid,  temperature,  and  the  presence  of 
resulting  compounds ;  one  product  is  usually  a  soluble  nitrate, 
though  some  metals,  such  as  tin  and  antimony,  form  in- 
soluble oxides.  The  gaseous  products  are  usually  oxides 
of  nitrogen,  especially  colorless  nitric  oxide  (NO),  which 
quickly  forms  brown  nitrogen  dioxide  (NO2)  in  the  air.  Hy- 
drogen is  seldom  liberated  so  that  it  can  be  collected;  it 
generally  reduces  the  nitric  acid  to  nitric  oxide  (NO)  and 
water.  Ammonia  gas  (NH3)  and  even  nitrogen  are  some- 
times formed.  The  reaction  between  moderately  dilute 
nitric  acid  and  copper  is  typical  of  some  metals  and  may  be 
written  thus :  — 

3  Cu  -f  8  HN03   =  3  Cu(N03)2  +  2  NO   +  4  H2O 

Copper  Nitric  Copper  Nitric  Water 

Acid  Nitrate  Oxide 

This  equation  is  really  made  up  of  three  equations  :  — 

(1)  2  HNO3  =  3  0  +  2  NO  +  H2O 

(2)  3  Cu  +  3  0  =  3  CuO 

Copper 
Oxide 

(3)  3  CuO  +  6  HN03  =  3  Cu(NO3)2  +  3  H2O 


220  INORGANIC  CHEMISTRY 

Eliminating  3  O  from  (1)  and  (2)  and  3  CuO  from  (2)  and 
(3),  the  remaining  terms  make  up  the  complete  equation. 
In  the  case  of  zinc,  which  is  typical  of  other  metals,  the 
equations  are :  — 

(1)  3  Zn  +  6  HNO3  =  3  Zn(NO3)2  +  6  H 

(2)  6  H  +  2  HNO3  =  2  NO  +  4  H2O 

Eliminating  the  common  factor  (6  H),  the  complete  equation 
is:  — 

3  Zn  +  8  HNO3  =  3  Zn(N03)2  +  2  NO  +  4  H2O 

Nitrates,  as  a  rule,  are  very  soluble  in  water.  They  behave 
in  various  ways  when  heated.  Some,  like  sodium  and  potas- 
sium nitrates,  lose  oxygen  and  pass  into  nitrites ;  others,  like 
copper  nitrate,  form  an  oxide  of  the  metal,  an  oxide  of  nitro- 
gen, and  oxygen ;  and  one,  ammonium  nitrate,  decomposes 
into  water  and  nitrous  oxide  (N2O).  Since  many  nitrates, 
when  heated,  give  up  oxygen,  they  are  powerful  oxidizing 
agents.  Potassium  nitrate  dropped  on  hot  charcoal  burns 
the  charcoal  vigorously  and  rapidly.  This  kind  of  chemical 
action  is  called  deflagration.  Nitrates  have  numerous  uses, 
and  these,  as  well  as  their  special  properties,  will  be  treated 
under  their  respective  metals. 

The  Test  for  Nitrates  (and  of  course  for  nitric  acid)  is  as 
follows :  Add  to  the  nitric  acid  or  the  solution  of  the  nitrate 
an  equal  volume  of  concentrated  sulphuric  acid,  and  cool  the 
mixture.  Upon  the  cool  mixture  pour  carefully  a  cold,  di- 
lute, freshly  prepared  solution  of  ferrous  sulphate.  A  dark 
brown  layer  appears  where  the  two  liquids  meet,  owing 
to  the  formation  of  an  unstable  compound  which  has  ap- 
proximately the  composition  represented  by  3  FeSO4 .  2  NO. 
Owing  to  the  solubility  of  nitrates,  the  nitrate  ion  (NO3~) 
cannot  be  precipitated. 


COMPOUNDS  OF  NITROGEN  221 

Nitrous  Acid,  HN02,  is  not  easily  obtained  in  the  free  state 
owing  to  its  instability,  but  the  nitrites  are  stable  compounds. 
Potassium  nitrite  (KNO2)  and  sodium  nitrite  (NaNO2)  are 
formed  by  removing  the  oxygen  from  the  corresponding 
nitrate  by  heating  with  lead.  Nitrites  give  a  brown  mix- 
ture of  nitric  oxide  and  nitrogen  dioxide  when  treated  with 
sulphuric  acid,  and  are  thus  readily  distinguished  from 
nitrates. 

OXIDES  OF  NITROGEN 
There  are  five  oxides  of  nitrogen  :  — 


NAME 

FORMULA 

Nitrous  Oxide              

N2O 

Nitric  Oxide            ....          .     .               

NO 

Nitrogen  Trioxide  .                    .                         . 

NoOq 

Nitrogen  Dioxide 

NO2 

N20B 

Only  three  of  these  are  important ;  viz.  nitrous  oxide,  nitric 
oxide,  and  nitrogen  dioxide. 

Nitrous  Oxide,  N2O,  is  one  of  the  numerous  decomposition 
products  of  nitric  acid,  but  it  is  usually  prepared  by  decom- 
posing ammonium  nitrate.  This  salt,  if  gently  heated,  first 
melts  and  then  decomposes  into  nitrous  oxide  and  water; 
the  gas  can  be  collected  over  water,  preferably  warm  water. 
The  equation  for  the  chemical  change  is  — 

NH4NO3  N2O         +        2H2O 

Ammonium  Nitrous  Water 

Nitrate  Oxide 

This  colorless  gas  has  a  sweet  taste  and  a  faint  but  pleasant 
odor.     It  is  less  soluble  in  hot  than  in  cold  water.     The 


222  INORGANIC   CHEMISTRY 

gas  does  not  burn,  but  it  supports  the  combustion  of  many 
burning  substances,  though  not  so  vigorously  as  oxygen  does. 
Sulphur,  for  example,  will  not  burn  in  nitrous  oxide,  unless 
the  sulphur  is  hot  and  well  ignited  at  first ;  very  fine  iron 
wire,  if  well  ignited,  burns  in  the  gas,  but  the  combustion  is 
not  so  conspicuous  as  in  oxygen.  The  products  of  the  chem- 
ical change  are  oxides  and  nitrogen.  The  most  striking 
property  of  nitrous  oxide  is  its  effect  on  the  human  system. 
If  breathed  for  a  short  time,  it  causes  more  or  less  nervous 
excitement,  sometimes  manifested  by  laughter.  The  gas 
was  called  "laughing  gas"  by  Davy.  If  breathed  in  large 
quantities,  it  slowly  produces  unconsciousness.  The  gas  is 
often  administered  when  unconsciousness  is  desired  for  a 
short  time,  as  in  dentistry.  It  is  easily  liquefied  by  cold,  by 
pressure,  or  by  botH  together,  since  its  critical  temperature 
is  about  +  37°  C.  It  is  often  used  in  liquid  form  to  furnish 
the  gas  itself  and  to  produce  low  temperatures. 

Nitrous  oxide  was  discovered  by  Priestley  in  1776;  but  its  com- 
position was  not  explained  until  1799,  when  Davy,  by  an  extensive 
study  of  its  properties,  proved  it  to  be  an  oxide  of  nitrogen.  In  his 
enthusiasm  Davy  wrote  a  friend:  "  This  gas  raised  my  pulse  upward  of 
twenty  strokes,  made  me  dance  about  the  laboratory  as  a  madman, 
and  has  kept  my  spirits  in  a  glow  ever  since."  It  is  needless  to  say 
that  the  usual  results  are  more  quieting. 

Nitric  Oxide,  NO,  has  long  been  known,  since  it  is  the 
usual  gaseous  product  formed  by  the  interaction  of  nitric 
acid  and  metals.  It  is  conveniently  prepared  by  the  inter- 
action of  copper  and  dilute  nitric  acid  (sp.  gr.  1.2).  The 
equation  for  the  complex  chemical  change  is  written  thus  :  — 

3Cu    +     8HN03     =       2  NO    +  3  Cu(NO3)2  +    4H20 

Copper  Nitric  Nitric  Copper 

Acid  Oxide  Nitrate 

The  gas  thus  prepared  is  impure,  and  it  is  customary  to  use 
ferrous  sulphate  and  nitric  acid  as  a  source  of  the  pure  gas. 


COMPOUNDS  OF  NITROGEN  223 

Nitric  oxide  is  a  colorless  gas,  but  upon  exposure  to  the  air 
it  combines  at  once  with  oxygen,  forming  dense  reddish 
brown  fumes  of  nitrogen  dioxide.  The  equation  for  this 
change  is  — 

2  NO  +  O2  =  2  NO2 

Nitric  Nitrogen 

Oxide  Dioxide 

This  property  distinguishes  nitric  oxide  from  all  other  gases. 
It  does  not  burn,  nor  does  it  support  combustion  unless  the 
burning  substance  (e.g.  phosphorus  or  sodium)  introduced 
is  hot  enough  to  decompose  the  gas  into  nitrogen  and  oxy- 
gen, and  then,  of  course,  the  liberated  oxygen  supports  the 
combustion.  It  is  only  slightly  soluble  in  water. 

Nitrogen  Dioxide  (or  Peroxide),  NO2,  is  the  reddish  gas 
formed  by  the  union  of  nitric  oxide  and  oxygen.  Thus :  — 

2  NO         +         O2  2N02 

Nitric  Nitrogen 

Oxide  Dioxide 

It  is  also  produced  by  heating  certain  nitrates.     Thus  :  — 
Pb(NO3)2  2N02         +         PbO         +         O 

Lead  Nitrogen  Lead  Oxygen 

Nitrate  Dioxide  Oxide 

The  fumes  of  nitrogen  dioxide  usually  appear  when  nitric 
acid  and  metals  interact,  but,  as  already  stated,  the  fumes 
are  not  produced  at  first,  being  the  result  of  a  second  chemical 
change  when  the  nitric  oxide  combines  with  oxygen  of  the 
air. 

Nitrogen  dioxide  has  a  disagreeable  odor,  and  if  breathed 
in  moderately  large  quantities,  it  is  poisonous.  It  interacts 
with  water  and  under  ordinary  conditions  yields  nitric  acid 
and  nitric  oxide.  The  gas  also  dissolves  in  concentrated 
nitric  acid,  forming  fuming  nitric  acid,  which  is  a  powerful 
oxidizing  agent. 


224  INORGANIC  CHEMISTRY 

When  the  gas  called  nitrogen  dioxide  is  sealed  in  a  tube  and 
cooled,  the  color  changes  from  red-brown  to  pale  yellow ;  a  yellow 
liquid  forms  at  about  26°  C.  and  a  nearly  colorless  solid  at  about 
-  12°  C.  The  yellow  substance  is  nitrogen  tetroxide  (N2O4). 
Upon  heating  the  tube  the  red-brown  gas  reappears;  at  about 
150°  C.  it  is  dark  red-brown  and  is  nitrogen  dioxide  (NO2).  These 
changes  show  that  at  ordinary  temperatures  the  tube  contains  a 
mixture  of  the  two  gases  which  are  in  equilibrium,  thus :  — 

N204     ^±     2  NO2 
Nitrogen  Nitrogen 

Tetroxide  Dioxide 

Although  the  red-brown  gas  as  ordinarily  seen  is  a  mixture,  it  is 
called  nitrogen  dioxide.  The  formation  of  nitrogen  dioxide  in- 
volves absorption  of  heat.  Hence  increased  temperature  favors 
the  forward  reaction  (see  Le  Chatelier's  law). 

Nitrogen  Trioxide,  N2O3,  and  Nitrogen  Pentoxide,  N2O6,  are  unstable 
compounds  and  have  no  practical  importance.  They  are  the  anhy- 
drides of  nitrous  and  nitric  acids. 

Aqua  Regia  is  an  old  term  used  by  the  alchemists  and  still 
applied  to  a  mixture  of  concentrated  nitric  and  hydrochloric 
acids  (1  vol.  to  3  vol.).  The  expression  means  "  royal  water," 
and  indicates  that  the  mixture  dissolves  gold  (a  "  noble  " 
metal),  which  is  insoluble  in  either  acid  alone.  Its  solvent 
power  depends  mainly  upon  the  free  chlorine  which  is  pro- 
duced in  the  mixture  by  the  oxidizing  action  of  the  nitric 
acid,  thus :  — 

HNO3  +  3  HC1  =  2  Cl  +  NOC1  +  2  H2O 

Aqua  regia  reacts  energetically  with  metals,  and  the  product 
of  the  reaction  is  always  a  chloride  of  the  metal. 

GAY-LUSSAC'S  LAW  OF  GAS  VOLUMES 

Several  gaseous  compounds  of  nitrogen  illustrate  Gay- 
Lussac's  law  of  gas  volumes.  Experiment  shows  the  follow- 
ing facts  about  these  gases  and  certain  others  previously 
studied :  — 


COMPOUNDS   OF  NITROGEN  225 

COMBINATION  OF  GASES  BY  VOLUME 


VOLUMES  OF  COMBINING  GASES 


VOLUMES  OF  GASEOUS  PRODUCTS 


2  vol.  hydrogen 
1  vol.  oxygen 

1  vol.  hydrogen 
1  vol.  chlorine 

3  vol.  hydrogen 

1  vol.  nitrogen 

2  vol.  nitrogen 
1  vol.  oxygen 

1  vol.  nitrogen 
1  vol.  oxygen 

1  vol.  nitrogen 

2  vol.  oxygen 

2  vol.  nitrogen 

3  vol.  oxygen 


2  vol.  water  vapor 

2  vol.  hydrochloric  acid  gas 

2  vol.  ammonia  gas 

2  vol.  nitrous  oxide  gas 

2  vol.  nitric  oxide  gas 

2  vol.  nitrogen  dioxide  gas 

2  vol.  nitrogen  trioxide  gas 


It  is  clear  from  the  above  table  that  small  whole  numbers 
express  the  relation  existing  between  the  volumes  of  the  com- 
bining gases  and  the  volume  of  the  gaseous  product.  This 
simple  relation  is  general  and  was  summarized  in  1808  by 
the  French  chemist  Gay-Lussac  in  the  form  of  a  law,  thus  :  — 

Gases  combine  in  volumes  which  bear  a  simple  numerical 
ratio  to  each  other  and  to  the  volume  of  their  gaseous  product. 

Additional  illustrations  of  this  fundamental  law  will  be 
given  in  subsequent  chapters.  (See  especially  Chapters  XV 
and  XVI  (oxides  of  carbon  and  hydrocarbons.)  See  also 
gas  equations,  Chapter  XIV.) 

PROBLEMS 

1.  How  many  grams  of  ammonia  gas  can  be  obtained  from  2140 
gm.  of  ammonium  chloride  by  heating  with  lime  ? 

2.  Calculate  the  percentage  composition  of   (a)  ammonium  chlo- 


226  INORGANIC  CHEMISTRY 

ride,  (6)  ammonium  hydroxide,  (c)  ammonium  sulphate,  (d)  am- 
monium nitrate. 

3.  What  weight  of  pure  sodium  nitrate  is  needed  to  produce  a 
metric  ton  of  pure  nitric  acid  ? 

4.  How  many  grams  of  each  element  in  27  gm.  of  pure  nitric 
acid? 

5.  What  weight  of  pure  nitric  acid  can  be  obtained  from  a  metric 
ton  of  sodium  nitrate  (95  per  cent  pure)  ? 

6.  Will  sodium  nitrate  or  potassium  nitrate  yield  the  greater 
weight  of  nitric  acid? 

7.  What  volume  of  sulphuric  acid  solution  having  a  specific 
gravity  of  1.8354  and  containing  93.19  per  cent  of  H2SO4  is  needed 
to  convert  10  kg.  of  pure  sodium  nitrate  into  pure  nitric  acid  ? 

8.  What  volume  of  nitric  acid  solution  having  a  specific  gravity 
of  1.4  and  containing  65.3  per  cent  of  HNO3  can  be  obtained  by  the 
interaction  of  sulphuric  acid  having  the  concentration  given  in 
Problem  7  and  25  metric  tons  of  pure  sodium  nitrate? 

9.  What  volume   of  oxygen   is  needed   to  combine  with  the 
hydrogen  obtained  by  passing  electric  sparks  through  150  cc.  of 
ammonia  gas  until  equilibrium  is  reached? 

10.  What  weight  and  what  volume  of   nitrous  oxide  can  be 
prepared  from  72  gm.  of  ammonium  nitrate  ?    (Standard  conditions. ) 

11.  What  weight  of  nitric  oxide  is  formed  by  the  interaction  of 
nitric  acid  and  45  gm.  of  copper?     What  weight  of  nitrogen  dioxide 
will  the  nitric  oxide  form? 

12.  What  weight  and  what  volume  of  oxygen  will  be  needed  to 
convert  (a)  70  gm.  and  (6)  70  1.  of  nitric  oxide  into  nitrogen  dioxide? 
(Standard  conditions.) 

13.  (a)  What  volume  of  oxygen  is  needed  to  convert  10  1.  of 
NO  into  NO2?     (b)  What  volume  of  nitrogen,  to  convert  12  1.  of 
hydrogen  into  NH3? 

14.  Calculate  the  volume  of  gas  formed  in  each  of  the  following 
reactions :    (a)  150  cc.  of  hydrogen  and  sufficient  nitrogen ;    (6)  hy- 
drogen and  150  cc.  of  nitrogen ;    (c)  100  cc.  of  oxygen  and  sufficient 
nitric  oxide. 

15.  Calculate  the  formulas  corresponding  to  (a)  N  =  46.666, 
O  =  53.333;  (6)N  =  40,  O  =  45.71,  H  =  14.28;  (c)N  =  35,  O  =  60, 
H  =  5.     What  is  the  name  of  each  compound  ? 


CHAPTER  XIV 
Atomic  and  Molecular  Weights  —  Valence 

EXTENDED  application  of  atomic  and  molecular  weights  has 
been  made  in  the  foregoing  pages.  In  the  present  chapter 
we  shall  consider  the  methods  by  which  both  atomic  and 
molecular  weights  are  determined;  several  cognate  prin- 
ciples of  fundamental  importance  will  also  be  discussed. 

Retrospect.  —  Before  beginning  the  development  of  atomic 
and  molecular  weights  it  will  be  helpful  to  review  certain  facts 
and  assumptions  which  bear  directly  upon  this  subject.  It 
will  be  recalled  that  chemical  compounds  have  a  definite  com- 
position by  weight,  and  furthermore  that  if  the  proportions 
of  the  elements  in  a  series  of  compounds  containing  the  same 
elements  are  expressed  in  a  special  way,  the  composition  is 
revealed  as  a  simple  multiple  relation.  Again,  it  will  be  re- 
membered that  gases  exhibit  a  striking  similarity  of  behav- 
ior not  only  when  they  are  subjected  to  heat  and  pressure 
but  also  when  they  undergo  chemical  transformations. 
These  phenomena  are  summarized  in  the  laws  of  Boyle, 
Charles,  and  Gay-Lussac,  which  have  already  been  discussed. 
It  has  also  been  seen  that  in  chemical  changes  neither  energy 
nor  matter  is  lost;  transformations  occur,  but  both  energy 
and  matter  are  conserved.  Finally  we  should  not  forget  that 
the  atomic  theory  offers  an  acceptable  explanation  of  the  com- 
position of  matter  and  of  certain  aspects  of  chemical  change 
by  assuming  that  atoms  are  the  gravimetric  units  of  chemical 
change  and  by  their  combinations  form  molecules  which  in 
turn  decompose  wholly  or  in  part,  thereby  producing  the 
varied  and  complicated  phenomena  succinctly  called  chem- 
ical changes.  These  atoms  have  a  relative  and  unvarying 

227 


228  INORGANIC   CHEMISTRY 

weight  called  the  atomic  weight  —  alike  for  each  atom  of  the 
same  element,  unlike  for  each  atom  of  different  elements.  So 
also  each  molecule  of  a  compound  has  a  relative  and  unvary- 
ing weight  called  the  molecular  weight,  which  is  the  sum  of 
the  weights  of  the  atoms  in  one  molecule  of  the  compound. 

Our  present  task  is  to  discuss  the  methods  of  determining 
these  relative  weights  of  atoms  and  molecules,  and  to  sup- 
plement the  data  collected  in  this  brief  retrospect  by  certain 
laws,  principles,  and  theories. 

Determination  of  Atomic  Weights.  —  The  atomic  weight 
of  an  element,  as  already  stated,  is  a  relative  weight.  It  is 
a  number  expressing  the  relation  of  the  weight  of  the  atom 
of  a  given  element  to  the  weight  of  the  atom  of  some  element 
chosen  as  a  standard.  Thus,  if  we  say  the  atomic  weight  of 
nitrogen  is  14,  we  mean  that  the  ratio  of  the  weight  of  the 
nitrogen  atom  and  the  weight  of  the  hydrogen  atom  is  14  to 
1,  provided  we  adopt  1  as  the  weight  of  the  hydrogen  atom; 
or  we  mean  that  the  ratio  of  the  weight  of  the  nitrogen  atom 
and  the  weight  of  the  oxygen  atom  is  14  to  16,  provided  we 
adopt  16  as  the  weight  of  the  oxygen  atom.  Hydrogen  was 
the  standard  for  many  years.  However,  since  oxygen  com- 
bines readily  with  a  large  number  of  elements  and  forms 
compounds  which  are  quite  suitable  for  experimental  work, 
the  oxygen  atom  has  been  adopted  as  the  international 
standard  atom  and  the  weight  16  has  been  given  to  it. 

When  compounds  are  analyzed,  the  results  show  the  pro- 
portions by  weight  in  which  the  constituent  elements  are 
combined.  If  one  molecule  of  a  compound  contained  only 
one  atom  each  of  the  united  elements,  the  relative  weights 
of  the  atoms  could  easily  be  determined.  Thus,  approxi- 
mately 8  parts  of  oxygen  combine  with  1  part  by  weight  of 
hydrogen  to  form  9  parts  of  water;  assuming  that  a  mole- 
cule of  water  contains  only  1  atom  of  each  element,  it  is 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE     229 

evident  .that  the  atomic  weight  of  hydrogen  would  be  2  if 
oxygen  is  16.  But  before  accepting  this  conclusion  we 
should  not  overlook  the  existence  of  another  compound  of 
hydrogen  and  oxygen  (called  hydrogen  peroxide)  in  which 
1  part  by  weight  of  hydrogen  combines  with  16  parts  of 
oxygen.  Assuming  as  before  that  a  molecule  of  hydrogen 
peroxide  contains  1  atom  of  each  element,  the  conclusion  is 
forced  upon  us  that  the  atomic  weight  of  hydrogen  would  be 
1  if  oxygen  is  16.  But  the  atomic  weight  of  hydrogen  can- 
not be  both  1  and  2 !  If  we  only  knew  which  of  the  com- 
pounds contained  a  single  atom  of  each  element  to  the 
molecule,  —  granting  for  the  time  being  that  such  is  the 
case,  —  the  problem  would  be  simple.  Analysis  does  not 
reveal  the  number  of  atoms  in  a  molecule.  Obviously  addi- 
tional data  are  indispensable. 

The  determination  and  final  selection  of  the  atomic  weight 
of  an  element  is  based  upon  (1)  the  equivalent  weight  of  the 
element,  (2)  the  molecular  weights  of  several  compounds  of 
the  element,  and  (3)  accurate  chemical  analysis  of  selected 
compounds  of  the  element. 

Equivalent  Weights.  —  The  chemical  analysis  of  a  com- 
pound, as  already  stated,  gives  the  proportion  by  weight 
in  which  the  constituents  are  combined.  If  these  propor- 
tions are  stated  in  a  special  way  instead  of  the  customary 
form  of  percentage,  certain  important  relations  are  revealed. 
Thus,  analysis  of  ferrous  oxide  yields  approximately  77.77  per 
cent  of  iron  and  22.22  per  cent  of  oxygen;  but  if  we  substitute 
8  for  22.22,  the  weight  of  iron  becomes  28,  because  22.22  and 
77.77  are  in  the  same  ratio  as  8  and  28.  When  a  similar 
process  of  modification  is  applied  to  the  percentage  of  the 
elements  in  different  compounds,  a  series  of  numbers  is  ob- 
tained known  as  the  equivalent  weights  of  the  elements. 
By  definition,  the  equivalent  weight  of  an  element  is  the 


230 


INORGANIC   CHEMISTRY 


number  of  grams  which  combines  with  or  displaces  8  gm.  of 
oxygen.  A  partial  summary  of  numerous  experiments  gives 
the  following :  — 

TABLE  OF  EQUIVALENT  WEIGHTS 


ELEMENT 

EQUIVALENT  WEIGHT 

ELEMENT 

EQUIVALENT  WEIGHT 

Oxygen    .... 

8 

Iron  (-ous)  . 

27.92 

Aluminium    .  . 

9.03 

Iron  (-ic)  .  . 

18.62 

Bromine  .... 

79.92 

Magnesium  . 

12.16 

Calcium  .... 

20.03 

Mercury  (-ic) 

100.3 

Carbon    .... 

3 

Potassium    . 

39.10 

Chlorine  .... 

35.46 

Silver  .... 

107.88 

Copper  (-ous)  . 

63.57 

Sodium  .  .   . 

23 

Copper  (-ic)  .  . 

31.79 

Sulphur.  .  . 

16.03 

Hydrogen  .  .  . 

1.008 

Zinc    .... 

32.68 

This  list  might  be  extended  to  include  all  the  elements 
which  form  compounds.  These  numbers  are  sometimes 
called  combining  numbers,  combining  weights,  or  simply 
equivalents.  The  term  equivalent  weights  is  preferable, 
because  they  actually  are  the  weights  chemically  equivalent 
to  each  other.  Thus,  if  we  start  with  hydrogen  chloride 
(HC1),  1  gm.  of  hydrogen  —  to  take  a  convenient  denomi- 
nation —  is  combined  with  35.46  gm.  of  chlorine,  and  this 
gram  of  hydrogen  can  be  replaced  chemically  by  32.68  gm. 
of  zinc,  12.16  gm.  of  magnesium,  39.10  gm.  of  potassium, 
23  gm.  of  sodium,  and  so  on.  These  elements  are  chem- 
ically equivalent  in  the  ratio  of  these  weights. 

Equivalent  weights  are  readily  found  by  experiment  in  most  cases. 
Various  methods  are  used,  and  the  equivalent  weight  is  not  always 
found  directly  in  terms  of  oxygen.  The  equivalent  of  hydrogen  is 
found  by  passing  hydrogen  over  hot  copper  oxide  —  as  in  the  deter- 
mination of  the  gravimetric  composition  of  water.  The  equivalent 
of  magnesium  is  found  by  filling  a  graduated  tube  with  dilute  hydro- 
chloric acid,  inserting  the  magnesium,  inverting  the  tube  in  a  dish, 
and  measuring  the  volume  of  the  liberated  hydrogen;  knowing  the 
weight  of  a  liter  of  hydrogen,  the  weight  of  the  liberated  hydrogen 


ATOMIC  AND   MOLECULAR  WEIGHTS  —  VALENCE     231 


can  be  calculated,  and  from  its  weight  the  weight  of  magnesium 
equivalent  to  1  gm.  of  hydrogen  can  be  found  by  proportion.  The 
equivalent  weight  of  magnesium  can  also  be  found  by  heating  a  known 
weight  of  magnesium  in  the  air  —  taking  care,  of  course,  not  to  lose 
the  product.  Equivalent  weights  of  copper  and  other  metals  as  well 
as  of  sulphur  and  of  carbon  can  be  found  by  passing  oxygen  over 
these  substances  in  a  tube  which  contains  or  is  attached  to  a  suitable 
apparatus  for  retaining  the  product.  The  interaction  of  metals  and 
acids  provides  a  simple  method  of  finding  the  equivalent  of  zinc, 
aluminium,  and  iron;  the  interaction  of  water  with  sodium  and 
with  calcium  permits  the  determination  of  the  equivalent  weights  of 
these  metals;  and  the  displacement  of  metals  from  solutions  of  their 
compounds  by  such  metals  as  zinc  and  magnesium  provides  another 
general  method. 

Equivalent  weights  have  been  very  carefully  determined 
by  experiment.  Comparison  of  the  equivalent  weight  with 
the  atomic  weight  of  the  same  element  reveals  an  important 
relation,  as  may  be  seen  by  the  following  :  — 

TABLE  OF  EQUIVALENT  WEIGHTS  AND  ATOMIC  WEIGHTS 


ELEMENT 

EQUIVALENT 
WEIGHT 

ATOMIC  WEIGHT 

MULTIPLE 

8 

16 

2 

A  lum  in  ium 

9  03 

27  1 

3 

Bromine       

79  92 

79.92 

1 

Calcium  . 

20  03 

40  06 

2 

Carbon    

3 

12 

4 

Chlorine 

35  46 

35  46 

1 

Copper  (-oils')  . 

63.57 

63.57 

1 

Copper  (-ic) 

31  79 

63  57 

2 

Hydrogen     '. 

1.008 

1.008 

1 

Iron  (-ous) 

27  92 

55  85 

2 

Iron  (-ic)     

18.62 

55.85 

3 

Magnesium                      . 

12  16 

24  32 

2 

100.3 

200.6 

2 

Potassium 

39  10 

39  10 

1 

Silver      

107.88 

107.88 

1 

Sodium 

23 

23 

1 

16.03 

32.07 

2 

32.68 

65  37 

2 

232  INORGANIC  CHEMISTRY 

An  examination  of  this  comparative  table  shows  an  integral 
relation  between  equivalent  weights  and  atomic  weights. 
In  other  words,  the  atomic  weight  of  an  element  is  identical 
with  its  equivalent  weight  or  is  a  simple  integral  multiple 
of  it.  The  significance  of  this  relation  will  be  discussed  after 
the  subject  of  molecular  weights  has  been  considered. 

Determination  of  Molecular  Weights.  —  Before  describing 
the  actual  methods  employed  in  determining  molecular 
weights,  it  will  be  necessary  to  discuss  the  kinetic  theory  of 
gases  and  Avogadro's  hypothesis.  These  theoretical  prin- 
ciples underlie  the  interpretation  of  the  results  obtained  by 
experiment  and  assist  in  the  correlation  of  the  properties  of 
gases'  summarized  by  the  laws  of  Boyle,  Charles,  and  Gay- 
Lussac. 

Kinetic  Theory  of  Gases  and  Avogadro's  Hypothesis.  - 
Extensive  study  of  gases  shows  that  in  general  they  conform 
to  fundamental  laws.  These  laws  indicate  the  uniform  and 
simple  structure  of  all  gases.  The  theory  proposed  to  ex- 
plain the  uniform  behavior  of  gases  is  called  the  kinetic 
theory.  According  to  this  theory  gases  are  conceived  to 
consist  of  molecules,  moving  constantly  and  rapidly  in  all 
directions;  these  particles  are  also  conceived  to  have  perfect 
elasticity,  and  in  their  movements  in  the  space  which  is  large 
compared  with  their  own  bulk  they  collide  with  each  other 
or  with  the  walls  of  the  containing  vessel,  rebound,  and  con- 
tinue to  move  without  loss  of  energy;  furthermore,  the  mole- 
cules of  a  gas  move  in  straight  lines  and  have  little  or  no 
tendency  to  repel  or  adhere  to  each  other;  i.e.  they  are  inde- 
pendent particles  separated  by  an  average  distance  much 
greater  than  their  own  diameters.  Recasting  this  theory 
into  a  concrete  form,  a  vessel  of  oxygen  gas  contains  a  vast 
number  of  molecules,  flying  about  rapidly  in  the  space, 
striking  each  other  and  the  walls  of  the  vessel,  rebounding 


ATOMIC  AND   MOLECULAR  WEIGHTS  —  VALENCE     233 

after  each  collision,  some,  eventually  all,  flying  out  and  min- 
gling with  the  air,  but  not  combining  with  each  other  or  com- 
pressing each  other,  or  even  adhering  (except  under  unusual 
conditions). 

Many  facts  about  gases  are  readily  interpreted  by  this 
theory.  For  example,  compressibility,  diffusion,  and  the 
ability  to  mingle  with  other  gases  find  explanation  in  the 
conception  of  the  constant  and  rapid  motion,  perfect  elas- 
ticity, and  relatively  great  separation  of  the  molecules. 
Boyle's  law  finds  explanation  in  the  conception  of  the  pres- 
sure produced  by  the  incessant  impacts  of  the  independent 
molecules  moving  in  straight  lines,  the  varying  pressure 
being  due  to  the  varying  number  of  impacts  in  a  given  time 
produced  in  the  total  space  within  which  the  gas  is  confined 
-  the  less  the  space,  the  greater  the  pressure.  Charles' 
law  is  explained  by  the  conception  of  the  change  in  the  veloc- 
ity of  the  molecules  due  to  a  change  in  temperature,  the 
uniformly  varying  volume  (at  a  constant  pressure)  being  due 
to  the  varying  frequency  (and  consequently  varying  energy) 
of  the  impacts  of  the  molecules  —  the  more  frequent  the 
number  of  impacts,  the  greater  the  pressure  in  a  constant 
volume,  or  what  is  the  same  thing,  the  greater  the  volume 
at  a  constant  pressure.  The  law  of  Gay-Lussac  (that  is, 
volumes  of  combining  gases  are  in  a  simple  integral  relation 
to  each  other  and  to  the  total  volume  of  the  gaseous  product) 
must  be  interpreted  jointly  by  the  atomic  theory  and  the 
kinetic  theory.  According  to  the  atomic  theory  chemical 
union  occurs  between  atoms,  while  according  to  the  kinetic 
theory  gases  consist  of  molecules.  Now,  when  gases  com- 
bine, is  the  combination  between  atoms  or  molecules?  Let 
us  attempt  to  answer  this  question  by  considering  the  forma- 
tion of  water  by  the  combination  of  oxygen  and  hydrogen 
gases.  According  to  Gay-Lussac's  law  two  volumes  of  hy- 
drogen and  one  volume  of  oxygen  produce  two  volumes  of 


234  INORGANIC  CHEMISTRY 

water  vapor.  The  simplest  assumption  (and  the  one  actually 
made  by  Dalton  when  this  question  was  first  asked)  is  that 
equal  volumes  of  elementary  gases  contain  the  same  number 
of  atoms.  Let  the  number  in  a  unit  volume  be  X.  Then 
in  our  illustrative  case  2  X  atoms  of  hydrogen  unite  with  X 
atoms  of  oxygen  to  form  2  X  particles  of  water  vapor ;  i.e. 
each  particle  of  water  vapor  contains  half  an  atom  of  oxygen ! 
But  according  to  the  atomic  theory  there  are  no  fractions 
of  atoms,  therefore  we  must  abandon  this  incorrect  assump- 
tion that  equal  volumes  of  gases  contain  an  equal  number  of 
atoms.  Another  assumption  can  be  made,  viz.  that  equal 
volumes  of  gases  contain  the  same  number  of  molecules. 
This  assumption  was  first  made  by  Avogadro,  an  Italian 
physicist,  and  is  still  known  and  accepted  in  a  slightly  modi- 
fied form  as  Avogadro's  hypothesis  (see  below).  If  this  as- 
sumption is  made,  then  the  facts  summarized  by  Gay-Lussac's 
law  can  be  satisfactorily  explained.  Let  the  number  of 
molecules  in  a  given  volume  of  a  gas  be  Z.  Applying  our 
assumption  to  the  previous  illustration,  2  Z  molecules  of 
hydrogen  combine  with  Z  molecules  of  oxygen  to  form  2  Z 
molecules  of  water  vapor.  Now  a  molecule  of  hydrogen  and 
of  oxygen  each  consists  of  at  least  two  atoms  (as  will  soon  be 
shown).  Therefore  the  illustration  can  be  interpreted  as 
follows :  The  given  volume  contains  Z  molecules.  The  2 
volumes  of  hydrogen  contain  2  Z  molecules  or  4  Z  atoms, 
while  the  1  volume  of  oxygen  contains  Z  molecules  or  2  Z 
atoms ;  or  6  Z  atoms  all  together.  When  these  volumes 
unite,  2  volumes  of  water  vapor  are  formed  which  contain  2  Z 
molecules  of  water  vapor.  Now  if  we  assume  (see  below) 
the  fact  that  a  molecule  of  water  vapor  consists  of  2  atoms 
of  hydrogen  and  1  of  oxygen,  it  is  obvious  that  not  only  are 
the  requisite  number  of  atoms  available  to  form  2  molecules 
of  water  vapor,  but  that  exactly  2  molecules  must  be  formed. 
In  a  few  words,  if  we  accept  the  view  that  gaseous  reactions 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE    235 

take  place  between  molecules,  then  Avogadro's  hypothesis 
follows  as  a  logical  conclusion. 

Avogadro's  hypothesis  is  usually  stated  thus  :  — 
There  is  the  same  number  of  molecules  in  equal  volumes  of 
all  gases  at  a  given  temperature  and  pressure. 

This  hypothesis  means,  for  example,  that  a  liter  of  hydrogen 
and  a  liter  of  oxygen  at  the  same  temperature  and  pressure 
contain  the  same  number  of  molecules.  It  therefore  fol- 
lows that  the  weights  of  the  molecules  in  equal  volumes  of 
these  two  gases  (and  all  other  gases)  at  a  given  temperature 
and  pressure  are  in  the  same  ratio  as  the  weights  of  their 
equal  volumes.  Hence  to  find  the  relative  weights  of  gaseous 
molecules,  it  is  only  necessary  to  determine  and  compare  the 
actual  weights  of  equal  volumes  of  the  gases  at  a  given  tem- 
perature and  pressure.  Furthermore,  if  we  express  the 
weights  of  molecules  in  terms  of  the  standard  adopted  for 
atomic  weights,  we  have  found  the  molecular  weight  of 
the  substance.  It  is  clear,  then,  that  the  molecular  weights 
of  elements  and  compounds  in  the  gaseous  state  can  be 
found  by  (1)  assuming  the  kinetic  theory  and  Avogadro's 
hypothesis,  and  (2)  determining  experimentally  the  rela- 
tive weights  of  gases. 

The  hypothesis  of  Avogadro  was  proposed  in  1811.  But  it  was 
not  favorably  received  nor  was  it  utilized  in  finding  molecular 
weights  until  about  1858.  At  this  time  it  was  shown  by  Canniz- 
zaro,  a  countryman  of  Avogadro,  to  be  a  reliable  hypothesis,  since 
its  application  yielded  molecular  weights  in  almost  complete  agree- 
ment with  the  weights  found  by  other  methods.  Or  more  strictly, 
the  atomic  weights  of  the  elements  derived  by  all  methods  (includ- 
ing the  method  which  involved  Avogadro's  hypothesis)  were  uni- 
form, accurate,  and  on  a  consistent  theoretical  basis.  Although  this 
hypothesis  cannot  be  verified  by  methods  usually  used  in  chemical 
investigations  and  does  not  have  the  certainty  of  a  demonstrable 
law,  it  is  in  harmony  with  the  laws  of  gases  and  can  be  logically 
deduced  from  the  kinetic  theory  of  gases. 


236  INORGANIC   CHEMISTRY 

Determination  of  Molecular  Weights  by  the  Vapor  Density 
Method.  —  The  density  of  a  gas  or  vapor  is  its  relative  weight, 
i.e.  its  weight  in  terms  of  a  standard.  Thus,  if  the  density  of 
hydrogen  is  accepted  as  1,  the  density  of  oxygen  is  about  16, 
because  a  given  volume  of  oxygen  weighs  about  sixteen  times 
as  much  as  an  equal  volume  of  hydrogen  at  the  same  tem- 
perature and  pressure.  Similarly,  if  the  density  of  air  is 
accepted  as  1,  the  density  of  oxygen  is  about  1.1  because  the 
weights  of  equal  volumes  under  the  same  conditions  are  in 
this  ratio.  We  may,  however,  choose  any  standard,  such  as 
hydrogen  =  2  or  oxygen  =  32,  since  the  different  values  can 
be  readily  transformed  into  each  other  when  we  know  the 
numerical  relation  of  the  standards. 

For  many  years  hydrogen  was  the  standard  gas  for  ex- 
pressing vapor  density,  and  the  molecular  weights  of  gases 
and  vapors  were  found  by  multiplying  their  vapor  densities 
by  2,  thus  :  — 

Molecular  Weight  =  Vapor  Density  referred  to  Hydrogen  X  2. 

The  vapor  density  is  multiplied  by  2  because  the  molecular 
weight  of  hydrogen  is  2;  and  the  molecular  weight  of  hydro- 
gen is  2  because  a  molecule  of  hydrogen  contains  at  least  two 
atoms  each  having  the  atomic  weight  1.  The  conclusion 
that  the  hydrogen  molecule  contains  at  least  two  atoms  is 
based  on  the  following  :  — 

One  volume  of  hydrogen  combines  with  one  volume  of 
chlorine  to  form  two  volumes  of  hydrogen  chloride  (HC1). 
Suppose  the  volume  of  hydrogen  contains  1000  molecules. 
Then,  according  to  Avogadro's  hypothesis,  the  equal  volume 
of  chlorine  will  contain  1000  molecules,  while  the  two  volumes 
of  the  product  will  contain  2000  molecules  of  hydrogen  chlo- 
ride. That  is :  — 

1000  molecules  of  Hydrogen  4- 1000  molecules  of  Chlorine 

=  2000  molecules  of  Hydrogen  Chloride. 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE     237 

Now  since  every  molecule  of  hydrogen  chloride  contains  at 
least  one  atom  each  of  hydrogen  and  chlorine,  the  2000 
molecules  must  contain  at  least  2000  atoms  each  of 
hydrogen  and  chlorine.  But  the  2000  atoms  of  hydrogen 
and  of  chlorine  were  provided  by  the  1000  molecules  of 
hydrogen  and  the  1000  molecules  of  chlorine.  Therefore 
each  molecule  of  hydrogen  and  of  chlorine  must  contain  at 
least  two  atoms.  There  is  no  evidence  which  leads  us  to 
believe  that  the  chlorine  or  the  hydrogen  molecule  contains 
more  than  two  atoms.  That  is,  no  case  is  known  in  which  a 
given  volume  of  hydrogen  gas  furnishes  the  material  for  more 
than  two  volumes  of  the  gaseous  product ;  the  same  is  true 
of  chlorine.  On  this  ground  we  base  our  belief  that  every 
hydrogen  and  every  chlorine  molecule  contains  only  two 
atoms. 

There  are  good  reasons,  however,  for  adopting  oxygen 
gas  =  32  as  the  standard  for  expressing  vapor  density.  The 
atomic  weight  of  oxygen  is  16,  and  this  is  the  international 
standard  for  atomic  weights.  Hence  if  we  express  molecular 
weights  in  terms  of  the  atomic  weight  of  oxygen,  these  values 
then  become  a  consistent  part  of  the  international  system  of 
expressing  the  quantitative  aspects  of  chemical  change. 
The  molecular  weight  of  oxygen  is  32,  because  a  molecule  of 
oxygen  contains  two  atoms  (each  weighing  16).  This  con- 
clusion is  based  on  the  following  :  — 

Two  volumes  of  hydrogen  unite  with  one  volume  of  oxygen 
to  form  two  volumes  of  water  vapor.  Suppose  a  single  vol- 
ume contains  1000  molecules.  Then  the  two  volumes  of 
water  vapor  must  contain  2000  molecules  and  each  mole- 
cule must  contain  at  least  one  atom  of  oxygen,  or  2000 
atoms  of  oxygen  in  all.  Since  2000  atoms  of  oxygen  were 
furnished  by  the  1000  molecules  of  oxygen,  each  molecule 
of  oxygen  must  contain  at  least  two  atoms;  and  no  facts 
lead  us  to  believe  that  there  are  more  than  two  atoms. 


238  INORGANIC  CHEMISTRY 

As  the  atomic  weight  of  oxygen  is  15,  the  molecular  weight 
is  32. 

A  method  of  determining  the  molecular  weight  of  a  gas 
(or  a  volatile  substance)  is  now  clear,  i.e.  multiply  the  vapor 
density  referred  to  oxygen  by  32. 

Another  method  is  used,  known  as  the  gram-molecular 
volume  method.  We  have  already  seen  (page  138)  that 
the  gram-molecular  weight  of  a  substance  is  the  number  of 
grams  numerically  equal  to  the  molecular  weight.  (The 
terms  molar  weight,  formula  weight,  and  mole  are  sometimes 
used  instead  of  gram-molecular  weight.)  E.g.  the  gram- 
molecular  weight  of  oxygen  is  32  gm.  Now  since  1  1.  of 
oxygen  weighs  1.429  gm.,  22.4  1.  approximately  (32  -f-  1.429 
=  22.393)  is  the  volume  of  oxygen  occupied  by  32  gm. 
According  to  Avogadro's  hypothesis  equal  volumes  of  all 
gases  at  the  same  temperature  and  pressure  contain  the 
same  number  of  molecules.  Hence  22.4  1.  of  any  gas  will 
have  a  weight  which  will  not  only  show  how  many  times 
heavier  the  gas  is  than  oxygen  but  which  will  also  be  the 
molecular  weight  of  the  gas  referred  to  oxygen  as  32.  This 
volume  (22.4  1.  at  0°  C.  and  760  mm.)  of  a  gas  is  called  its 
gram-molecular  volume.  Hence  the  second  method  of  deter- 
mining the  molecular  weight  of  a  gas  is  to  find  the  weight 
in  grams  of  22.4  1.  (at  0°  C.  and  760  mm.)  of  the  gas. 

Experimental  Determination  of  Molecular  Weights.  — 
The  weight  of  one  liter  of  oxygen  has  been  carefully  deter- 
mined. Hence  the  simplest  method  of  determining  the 
molecular  weight  of  a  gas  or  vapor  would  appear  to  be 
merely  to  find  the  exact  weight  of  a  liter  of  the  gaseous  or 
vaporized  substance  and  make  the  necessary  calculation. 
It  is  more  convenient  experimentally,  however,  to  find  the 
volume  of  air  displaced  by  the  vapor  of  a  known  weight  of 
a  substance  and  then  calculate  the  weight  of  22.4  1. 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE    239 


A  vapor  density  method  frequently  used  is  one  devised  by  Vic- 
tor Meyer.  A  simplified  form  of  the  apparatus  is  shown  in  Figure 
26.  The  bulb  B  of  the  inner  tube  is  heated  to" 
a  constant  temperature  by  the  vapor  of  the  liquid 
boiling  in  the  larger  tube  A.  The  gas  measur- 
ing tube  D,  filled  with  water,  is  then  inverted 
in  the  vessel  of  water  E  over  the  end  of  the 
capillary  tube  C.  Finally,  a  weighed  quantity 
of  the  substance  (in  a  small  bulb  or  bottle) 
is  introduced  into  B  by  quickly  removing  and 
replacing  the  stopper  F ;  a  wad  of  glass  wool  or 
asbestos  at  the  bottom  of  B  prevents  the  tube 
from  being  broken.  The  substance  soon  vapor- 
izes, and  the  vapor  forces  its  own  volume  of 
air  into  the  gas  tube  E.  When  the  substance 
is  completely  vaporized,  the  volume  of  air  in 
D  is  measured  and  reduced  to  the  volume  it 
would  occupy  if  it  were  a  dry  gas  at  0°  C. 
and  760  mm.  From  the  corrected  volume  and 
the  weight  of  the  substance  the  weight  of  22.4  1. 
is  calculated.  For  example,  .1008  gm.  of  chloro- 
form displaced  18.93  cc.  of  air  (corrected 
volume).  If  18.93  cc.  of  chloroform  vapor 
weigh  .1008  gm.,  22.4  1.  will  weigh  118.3  gm. 
That  is,  according  to  this  experiment,  the  molec- 
ular weight  of  chloroform  is  118.3  (the  exact 
value  being  119.5).  FIG.  26.  —  Appara- 

tus  for   determin- 

The  vapor  density  method  is  limited  to      ing  vapor  density- 
gaseous    or    volatile    substances.      Other 
methods  are  now  available,  viz.  the  freezing  point  and  boiling 
point  methods. 

In  Chapter  IX  it  was  stated  that  the  freezing  point  of  a 
solution  is  lower  than  the  freezing  point  of  the  solvent,  and 
that  the  depression  of  the  freezing  point  in  dilute  solutions 
is  approximately  proportional  to  the  concentration  of  the 
solution  in  the  case  of  all  substances  which  are  not  ionized  or 
do  not  unite  with  the  solvent.  Furthermore,  by  extending 
the  results  obtained  by  experiments  with  dilute  solutions,  it 


240 


INORGANIC  CHEMISTRY 


has  been  found  that  if  a  mole  (i.e.  a  number  of  grams  numer- 
ically equal  to  the  molecular  weight)  of  substances  which 
depress  the  freezing  point  normally  is  dissolved  in  1000  gm. 
of  water,  the  freezing  point  of  all  such  solutions  is  depressed 
the  same  number  of  degrees ;  i.e.  each  solution  freezes  at 
approximately—  1.86°  C.  This  means,  for  example,  that  a 
solution  of  342  gm.  of  sugar  (C^H^On)  in 
1000  gm.  of  water  freezes  at  approxi- 
mately —  1.86°  C.  Since  this  number 
(  —  1.86)  is  the  same  for  all  solutions 
containing  the  molecular  weight  in  grams 
in  1000  gm.  of  water,  it  is  sometimes 
called  the  molecular  depression  constant 
(K).  Now,  if  we  find  by  experiment  the 
amount  of  depression  of  the  freezing 
point  caused  by  a  solution  of  known 
concentration,  the  molecular  weight  of 
the  solute  is  readily  calculated.  An 
example  will  make  this  point  clear.  It 
was  found  that  a  solution  of  50  gm.  of 
methyl  alcohol  in  1000  gm.  of  water  froze 
at  approximately  —  2.90°  C.  A  solution 
containing  one  mole  (i.e.  a  gram-molec- 
ular weight)  of  methyl  alcohol  in  1000 
gm.  of  water  would  freeze  at  approxi- 
mately —  1.86°  C.  Hence  the  molecular 
weight  in  grams  is  found  by  the  pro- 
portion — 

50:  x::  -2.90  :  -  1.86 

or  x  =  32.     That  is,  32  is  the  molecular 
weight  of  methyl  alcohol. 


FIG.  27.  —  Beckmann 
apparatus  for  finding 
molecular  weights  by 
depression  of  the 
freezing  point. 


This  method  of  finding  molecular  weights  was  first  applied  ex- 
tensively by  Raoult  (1830-1901)  and  is  sometimes  called  the  cryo- 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE    241 

scopic  method.  The  experiment  may  be  performed  in  an  apparatus 
devised  by  Beckmann  (Fig.  27).  The  solvent  is  placed  in  the  inner 
tube  A  provided  with  a  side  tube  E  through  which  the  solute  is 
introduced.  A  very  accurate  and  sensitive  thermometer  T  passes 
through  a  stopper  into  the  inner  tube,  which  is  also  provided  with 
a  stirrer  S.  An  outer  tube  C  serves  as  an  air  jacket,  and  is  sur- 
rounded by  the  freezing  mixture,  which  is  placed  in  the  large  vessel 
B.  A  weighed  quantity  of  the  solvent  is  put  into  the  tube  A,  and 
its  freezing  point  is  carefully  determined.  Then  a  weighed  quan- 
tity of  the  solute  is  introduced  through  the  side  tube  E,  and  the 
freezing  point  of  the  solution  is  determined.  The  difference  be- 
tween the  two  freezing  points  is  the  depression,  and  from  this  de- 
pression the  molecular  weight  can  be  calculated.  Let  us  take  an 
example.  Suppose  4.98  gm.  of  sugar  (dsH^On)  are  dissolved  in 
96.9  gm.  of  water  and  the  depression  is  .287°  C.  Since  96.9  gm.  of 
water  contain  4.98  gm.  of  sugar,  1000  gm.  of  water  would  contain 
51.39  gm.  of  sugar.  If  51.39  gm.  of  sugar  cause  a  depression  of 
.287°,  then  the  number  of  grams  which  would  cause  the  molecular 
depression  can  be  found  by  the  proportion  — 

51.3^:. 287::  x:  1.86 
or  x  =  333.     The  correct  value  is  342. 

Molecular  weights  can  also  be  determined  by  an  analogous 
method  known  as  the  boiling-point  method. 

Exact  and  Approximate  Molecular  Weights.  — The  mo- 
lecular weights  found  by  experiment  are  only  approximate ; 
that  is  they  are  not  exactly,  though  often  very  nearly,  equal 
to  the  sum  of  the  exact  atomic  weights  in  one  mole- 
cule. The  difference  is  due  partly  to  the  difficulties  in 
making  accurate  measurements  of  temperature,  partly  to  im- 
purities in  the  substances  and  slight  defects  in  the  experi- 
mental method,  and  partly  also  to  the  erroneous  assumption 
that  Avogadro's  hypothesis  is  absolutely  true  for  all  tempera- 
tures. But  these  errors,  however,  do  not  affect  the  validity 
of  the  general  result.  Most  molecular  weights  used  in  chemi- 
cal discussions  and  calculations  are  calculated  molecular 
weights.  That  is,  after  the  composition  of  the  substance  has 


242  INORGANIC   CHEMISTRY 

been  determined  by  analysis,  the  molecular  weight  found  by 
experiment  is  slightly  changed  so  that  it  will  equal  the  sum 
of  the  weights  of  the  atoms  in  a  single  molecule.  It  should 
be  noted  that  the  failure  to  find  exact  molecular  weights  by 
experiment  is  not  a  serious  misfortune.  The  goal  is  atomic 
weights,  and  if  these  are  accurately  determined  (as  they  can 
be),  the  exact  molecular  weight  is  readily  found  by  merely 
adding  the  atomic  weights  corresponding  to  the  number  of 
atoms  in  the  molecule.  But,  as  repeatedly  stated,  approxi- 
mate molecular  weights  must  be  known  in  order  to  find  the 
atomic  weights  of  the  elements.  It  should  be  noted  further 
that  the  methods  of  finding  molecular  weights  described  in 
the  foregoing  pages  apply  only  to  substances  in  the  gaseous 
or  dissolved  state.  No  experimental  method  is  known  for 
determining  the  molecular  weights  of  substances  in  the 
solid  state.  Therefore,  unless  there  is  evidence  to  the  con- 
trary, the  molecular  weight  of  a  substance  in  the  solid  state 
is  assumed  to  have  the  same  value  as  that  found  by  the  usual 
methods. 

Relation  of  Atomic  Weights  to  Molecular  Weights  and 
Equivalent  Weights  and  Determination  of  Atomic  Weights. 
—  It  has  been  shown  that  the  atomic  weights  of  the  elements 
bear  a  simple  numerical  relation  to  molecular  weights  and 
equivalent  weights.  Furthermore,  the  methods  of  deter- 
mining and  calculating  both  molecular  and  equivalent  weights 
have  been  described.  It  is  now  appropriate  to  discuss  the 
methods  of  selecting  that  weight  known  as  the  atomic 
weight,  which  bears  the  correct  relation  to  the  equivalent 
weight  on  the  one  hand  and  the  molecular  weight  on  the 
other. 

The  first  method  we  shall  consider  might  be  called  the 
minimum  weight  method.  It  consists  in  (a)  determining  by 
appropriate  chemical  and  physical  methods  the  molecular 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE    243 

weights  of  several  representative  compounds  of  the  element, 
then  (b)  finding  the  weight  of  the  element  contained  in  the 
molecular  weight  of  each  compound,  and  finally  (c)  select- 
ing the  minimum  value  from  the  values  obtained  in  (b). 
When  this  process  is  applied  to  several  of  the  well  known 
elements,  the  results  may  be  tabulated  (in  round  numbers) 
as  follows :  — 


65 

to 

to 

w 

B 

fa 

M 

&  S 

M 

H 

rt 

b 

COMPOUND 

o  2 

8 

O 

O 

H 

E 

• 

^ 
o 

p 
P 
H 

| 

M 

fe 

0 

CO 

Water 

]8 

16 

2 

Hydrogen  Dioxide     . 

34 

32 

2 









Hydrogen  Chloride  .     . 

36.5 

— 

1 

35.5 

— 

— 

— 

Ammonia    .... 

17 

3 

14 

Nitric  Acid 

63 

48 

1 

14 

Nitrous  Oxide  .... 

44 

16 



28 





Nitric  Oxide     .... 

30 

16 

— 

— 

14 

— 

— 

Nitrogen  Dioxide  .     .     . 

46 

32 

— 

— 

14 

— 

— 

Carbon  Monoxide      .     . 

28 

16 

— 

— 

— 

12 

— 

Carbon  Dioxide    .     . 

44 

32 

— 

— 

— 

12 



Methane 

16 



4 

12 

Ethylene 

28 

4 

24 

Acetylene         .     . 

26 

2 

24 

Ether 

74 

16 

10 

48 

Ethyl  Alcohol  .... 

46 

16 

6 

— 

— 

24 



Chloroform      .... 

119.5 

— 

1 

106.5 

— 

12 

— 

Carbon  Tetrachloride    . 

154 

— 



142 



12 



Cyanogen  Chloride    . 

61.5 

— 

— 

35.5 

14 

12 

— 

Sulphur  Dioxide  .     .     . 

64 

32 

— 

— 

— 

— 

32 

Sulphur  Trioxide  .     .     . 

80 

48 

— 

— 

— 

— 

32 

Carbon  Disulphide    .     . 

72 

— 

— 

— 

— 

12 

64 

Hydrogen  Sulphide  .     . 

34 

— 

2 

— 

— 

— 

32 

Sulphuryl  Chloride   .     . 

135 

32 

— 

71 

— 

— 

64 

Minimum  weight  of  each  element 

16 

1 

35.5 

14 

12 

~w 

244  INORGANIC  CHEMISTRY 

In  this  table  columns  one  and  two  contain  the  names  of 
the  compounds  and  their  approximate  molecular  weights; 
the  other  columns  contain  the  parts  of  the  molecular  weights 
that  belong  to  the  atoms  of  the  elements  in  a  molecule  of 
the  compound.  For  example,  the  procedure  in  the  case  of 
water  is  as  follows :  (1)  By  experiment  the  approximate 
molecular  weight  of  water  vapor  is  found  to  be  18 ;  (2)  by 
analysis  the  compound  is  shown  to  contain  88.82  per  cent 
oxygen  and  11.18  per  cent  hydrogen;  and  (3)  the  products 
of  18  and  these  percentages  are  16  and  2  respectively.  An 
examination  of  the  weights  of  the  elements  shows  that  in 
each  column  (1)  the  minimum  weight  is  O  =  16,  H  =  1, 
Cl  =  35.5,  N  =  14,  C  =  12,  and  S  =  32,  and  (2)  the  other 
weights  are  simple  multiples.  These  facts  are  significant. 
The  smallest  weights  must  be  the  weights  of  single  atoms, 
for  it  is  highly  probable  that  one  or  more  compounds  in  a 
representative  group  will  contain  only  one  atom  of  a  given 
element ;  and  in  these  compounds,  of  course,  the  part  of  the 
molecular  weight  apportioned  to  the  element  in  question  is 
the  atomic  weight.  In  the  other  compounds  that  contain 
this  element  the  part  of  the  molecular  weight  apportioned 
to  the  element  will  be  a  multiple  of  the  atomic  weight ; 
obviously  these  compounds  contain  two  or  more  atoms  of 
the  element.  Thus,  in  hydrogen  dioxide  the  weight  appor- 
tioned to  oxygen  is  twice  that  in  water,  and  hence  a  mole- 
cule of  hydrogen  dioxide  contains  two  atoms  of  oxygen. 

A  comparison  of  the  atomic  weights  and  the  equivalent 
weights  of  the  elements  shows  that  the  atomic  weight  of 
an  element  is  equal  to  the  equivalent  weight  or  to  a  small 
integral  multiple  of  it.  (See  table  on  p.  231.)  This  rela- 
tion now  demands  a  fuller  explanation  than  hitherto  given. 
The  equivalent  weight,  as  previously  stated,  is  a  number 
obtained  by  expressing  in  a  special  way  the  per  cent  in 
which  elements  combine  or  exchange  places.  It  is  simply 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE    2i5 

the  number  of  parts  by  weight  of  an  element  which  com- 
bine with  or  displace  8  parts  by  weight  of  oxygen.  But 
the  equivalent  weight  is  not  the  unit  used  in  chemistry  to 
express  the  composition  of  compounds  or  the  quantitative 
relations  of  the  elements.  The  atomic  weight  is  the  quan- 
titative unit.  Hence  the  equivalent  weights  must  be  multi- 
plied by  certain  integers  in  order  to  transform  the  equivalent 
weight  of  an  element  into  the  corresponding  atomic  weight. 
That  is  to  say,  the  equivalent  weights  are  the  fundamental 
empirical  proportions,  while  the  atomic  weights  are  the  ad- 
justed gravimetric  units  deduced  from  the  equivalent  weights 
on  the  one  hand  and  from  the  molecular  weights  on  the 
other  hand. 

Another  method  for  determining  atomic  weights,  appli- 
cable to  the  solid  elements  only,  is  known  as  Dulong  and 
Petit's  method  because  it  utilizes  an  approximate  law  an- 
nounced and  applied  by  them  about  1819.  The  law,  com- 
monly called  the  law  of  specific  heats,  may  be  stated  as 
follows :  — 

The  product  of  the  specific  heat  and  atomic  weight  of  the  solid 
elements  is  approximately  equal  to  6.25. 

By  specific  heat  is  meant  the  quantity  of  heat  necessary  to 
raise  the  temperature  of  a  substance  one  degree  compared 
with  the  quantity  necessary  to  raise  the  temperature  of  the 
same  weight  of  water  one  degree.  If  the  same  quantity  of 
heat  is  imparted  to  equal  weights  of  water  and  mercury, 
the  temperature  of  the  mercury  will  be  much  higher  —  about 
32  times  higher  than  that  of  the  water.  That  is,  the  mercury 
requires  only  about  -fa  as  much  heat  as  the  water.  In  other 
words,  the  specific  heat  of  mercury  is  -fa,  or  .0312.  The 
specific  heat  of  elements  in  the  solid  state  can  be  found 
readily  by  experimental  methods.  The  number  found  by 
multiplying  the  specific  heat  of  a  solid  element  by  its  atomic 
weight  varies  somewhat,  but  in  many  cases  it  is  between 


246 


INORGANIC   CHEMISTRY 


6  and  6.5  (approximately  6.25).     This  relation  is  illustrated 
by  the  following  :  — 

TABLE  OF  SPECIFIC  HEATS 


ELEMENT 

SPECIFIC  HEAT 

APPROXIMATE 
ATOMIC  WEIGHT 

PRODUCT 

Calcium      

.170 

40 

6.8 

.095 

63.5 

6.03 

Iron  ...         .... 

.114 

56 

6.38 

.031 

207 

6.41 

Potassium  

.166 

39 

6.47 

.293 

23 

6.73 

Sulphur  

.178 

32 

5.7 

Tin     

.055 

119 

6.54 

Zinc  

.094 

65 

6.11 

Dulong  and  Petit's  law  is  only  an  approximation,  and  it 
serves  merely  to  check  values  obtained  by  other  methods 
and  to  show  whether  the  atomic  weight  of  an  element  is  a 
multiple  of  the  equivalent,  or  identical  with  it.  The  use  of 
the  law  in  checking  atomic  weights  may  be  illustrated  as 
follows :  The  specific  heat  of  silver  was  found  to  be  .057; 
if  6.25  is  divided  by  this  number,  the  quotient  is  about  109. 
This  result  shows  that  the  atomic  weight  of  silver  is  certainly 
not  55  or  218,  but  is  approximately  109  (the  exact  value 
being  107.88).  Again,  the  atomic  weight  of  uranium  may  be 
238  or  119  according  to  the  chemical  analysis;  but  only  the 
former  conforms  to  Dulong  and  Petit's  law,  and  hence  it  is 
accepted  as  the  approximate  atomic  weight. 

When  the  approximate  atomic  weight  of  an  element  has 
been  chosen  on  the  basis  of  the  foregoing  principles,  the  task 
still  remains  to  determine  the  accurate  value  of  this  impor- 
tant weight  by  chemical  analysis.  The  general  method  em- 
ployed can  be  illustrated  by  a  determination  made  by  the 
American  chemist  Richards,  whose  work  is  very  accurate. 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE    247 

He  found  that  28.26299  gm.  of  silver  chloride  were  formed 
from  21.27143  gm.  of  silver.  He  accepted  AgCl  as  the 
formula  of  silver  chloride  and  107.880  as  the  atomic  weight 
of  silver,  and  calculated  the  atomic  weight  of  chlorine 

thus : — 

28.26299  -  21.27143  =  6.99156 

Wt.  of  Silver :  Wt.  of  chlorine  : :  At.  wt.  of  silver  :  At.  wt.  of  chlorine 
21.27143     :        6.99156        ::         107.880  x 

x  =  35.458 

The  international  atomic  weight  of  chlorine  (35.46)  is 
based  on  this  and  other  determinations  made  by  the  same 
chemist. 

The  exact  determination  of  atomic  weights  is  a  difficult 
task.  And  although  the  utmost  care  is  used  in  purifying 
the  chemicals  and  performing  the  analysis,  the  results  of 
different  experimenters  do  not  always  exactly  agree.  There- 
fore an  international  committee  was  chosen  several  years 
ago  to  select  the  most  accurate  atomic  weights  of  the  ele- 
ments. These  weights  are  embodied  in  a  table  published 
annually  and  called  the  International  Table  of  Atomic 
Weights.  The  entire  table  is  given  in  the  Appendix,  §  5,  and 
a  supplementary  table  is  given  on  the  inside  of  the  back 
cover.  In  the  latter  table  the  accepted  atomic  weights 
are  placed  in  one  column  and  the  approximate  values  in 
another.  The  approximate  atomic  weights  are  often  suffi- 
ciently accurate  for  general  reference  and  in  making  chemical 
calculations;  they  may  be  used  in  solving  the  problems  in 
this  book. 

The  methods  and  principles  used  in  determining  the  atomic  weight 
of  an  element  can  be  reviewed  by  the  general  plan  which  would  be 
followed  out  in  the  case  of  zinc  (if  its  atomic  weight  were  unknown), 
(a)  The  equivalent  of  zinc  is  found  by  experiment  to  be  32.68. 
Therefore  the  atomic  weight  is  32.68,  or  some  multiple  of  this  num.- 


248  INORGANIC  CHEMISTRY 

her.  (6)  The  molecular  weight  of  zinc  chloride  is  found  by  its  vapor 
density  to  be  136.  (c)  Analysis  of  zinc  chloride  yields  47.8  per  cent 
of  zinc.  Therefore,  47.8  per  cent  of  136,  or  65.08,  is  zinc.  That  is, 
this  number  65.08  is  the  weight  of  the  smallest  part  of  zinc  in  zinc 
chloride,  and  it  may  be  the  weight  of  one  atom,  (d)  The  specific 
heat  of  zinc  is  about  .094.  Applying  Dulong  and  Petit's  law,  the 
approximate  atomic  weight  is  found  to  be  66.4  (i.e.  6.25  -=-  .094). 
Therefore,  the  atomic  weight  is  about  65,  and  not  32.68.  (e)  Careful 
analysis  of  zinc  compounds  shows  that  the  atomic  weight  of  zinc 
(on  the  basis  O  =  16)  is  65.37. 

Determination  of  Formulas  of  Compounds.  —  The  formula 
of  a  compound  is  an  expression  of  its  composition;  that  is, 
it  is  a  group  of  symbols  which  not  only  expresses  the  pro- 
portions of  the  weights  of  the  elements  in  a  compound,  but 
also  the  number  of  atoms  whose  sum  equals  the  molecular 
weight.  In  a  word,  a  formula  is  a  molecular  formula,  pro- 
vided, of  course,  the  molecular  weight  is  known.  For  ex- 
ample, the  proportion  of  hydrogen  to  chlorine  in  hydrogen 
chloride  is  1  to  35.5,  and  the  molecular  weight  is  known  to 
be  36.5;  therefore,  there  must  be  one  atom  each  of  hydrogen 
and  chlorine  in  a  molecule,  and  the  formula  is  HC1. 

Again,  the  proportion  of  carbon  to  hydrogen  in  a  certain 
hydrocarbon  is  12  to  2,  and  its  vapor  density  (referred  to 
oxygen)  is  5.09.  The  formula  would  be  CH2  according  to  the 
proportions  by  weight.  But  the  molecular  weight  of  such  a 
compound  would  be  only  14  (i.e.  12  +  2),  whereas  the  vapor 
density  of  the  original  hydrocarbon  necessitates  the  molec- 
ular weight  162.88  (i.e.  32  x  5.09),  which  is  nearly  twelve 
times  the  weight  corresponding  to  CH2.  Hence  the  molec- 
ular formula  is  not  CH2,  but  Ci2H24.  Again,  the  proportion 
of  hydrogen  to  oxygen  in  hydrogen  peroxide  is  1  to  16,  and 
the  formula  HO  would  express  this  proportion.  But  the 
molecular  weight  of  hydrogen  peroxide  has  been  found  by 
the  freezing  point  method  to  be  nearly  34.  Therefore, 
H2O2,  not  HO,  must  be  the  formula,  because  H2O2  corre- 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE      249 

spends  to  the  molecular  weight  as  well  as  to  the  proportions 
by  weight. 

If  the  composition  of  a  compound  is  known,  the  smallest 
number  of  atoms  corresponding  to  the  composition  can  be 
readily  calculated  by  dividing  the  per  cent  of  each  element 
in  the  compound  by  the  atomic  weight  (and,  if  necessary, 
reducing  the  quotients  to  "the  smallest  whole  numbers) .  That 
is,  we  merely  transform  the  percentage  into  numbers  which 
express  the  atomic  relations  by  distributing  the  propor- 
tional parts  among  the  elements  according  to  that  method 
of  expressing  composition  adopted  in  chemistry.  Some 
examples  will  make  this  point  clear.  According  to  analysis, 
a  compound  contains  40  per  cent  of  calcium,  12  per  cent  of 
carbon,  and  48  per  cent  of  oxygen.  Dividing  each  per  cent 
by  the  proper  atomic  weight,  we  have  :  — 

40  ^  40  =  1 
12  -5-12=1 
48-^16=3 

That  is,  one  molecule  of  this  compound  contains  (at  least) 
one  atom  each  of  calcium  and  carbon,  and  three  atoms  of 
oxygen.  Hence,  the  simplest  formula  is  CaCO3.  Again,  a 
compound  was  found  to  contain  2.04  per  cent  of  hydrogen, 
32.65  per  cent  of  sulphur,  and  65.31  per  cent  of  oxygen. 
Proceeding  as  in  the  previous  case,  we  have  :  — 

2.04-^1  =2.04 
32.65-^-32=1.02 
65.31 --16  =  4.08 

Reducing  these  quotients  to  the  smallest  set  of  integral 
numbers,  we  have  :  — 

2.04-^1.02  =  2 

1.02 -*- 1.02  =  1 

4.08-^1.02  =  4 


250  INORGANIC  CHEMISTRY 

That  is,  one  molecule  of  this  compound  contains  two  atoms 
of  hydrogen,  one  of  sulphur,  and  four  of  oxygen.  Hence 
the  simplest  formula  is  H2SO4.  If  the  molecular  weight  of 
a  compound  cannot  be  found,  then  the  simplest  formula 
(found  as  in  the  above  cases)  is  usually  accepted  as  the 
molecular  formula.  Finally,  a  compound  was  found  by 
analysis  to  contain  92.3  per  cent  of  carbon  and  7.7  per  cent 
of  hydrogen;  the  vapor  density  (referred  to  oxygen)  was 
2.4375.  Proceeding  as  above,  we  have  :  — 

92.3-^12  =  7.7  or  1 
7.7 -s-    1  =  7.7  or  1 

That  is,  the  compound  contains  at  least  one  atom  of  car- 
bori  and  hydrogen,  and  would  have  the  formula  CH,  if 
nothing  were  known  about  its  molecular  weight.  The  vapor 
density  2.4375  requires  the  molecular  weight  78,  which  is  six 
times  the  weight  (13)  corresponding  to  the  formula  CH. 
Hence  the  correct  formula  of  this  compound  is  not  CH 
but  C6H6. 

To  recapitulate :  The  simplest  formula  of  a  compound 
is  found  by  dividing  the  per  cent  of  each  element  by  its 
atomic  weight  and  reducing  these  quotients  to  the  smallest 
whole  numbers  (if  necessary) ;  if  the  molecular  weight  is  not 
known,  this  formula  is  accepted  as  the  molecular  formula. 
The  molecular  formula  of  a  compound  is  found  by  three 
steps :  (a)  Find  the  simplest  formula,  (b)  divide  the  molecular 
weight  by  the  sum  of  the  weights  of  the  atoms  in  the  simplest 
formula,  (c)  multiply  the  integral  numbers  of  the  simplest 
formula  by  the  quotient  obtained  in  (6).  (Compare  illustra- 
tion in  preceding  paragraph.) 

Molecular  Weights  and  Molecular  Formulas  of  Elements. 

—  Several  elements  are  gases  at  ordinary  temperatures  and 

others  can  be  vaporized  by  heating.     Hence  their  molecular 

weights  can  be  determined  by  the  vapor  density  method. 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE    251 

When  this  is  done,  the  results  indicate  that  the  elements 
fall  into  three  classes.  (1)  The  gaseous  elements  already 
studied,  as  well  as  some  others,  have  molecular  weights 
which  are  twice  the  atomic  weight ;  that  is,  the  molecule 
consists  of  two  atoms,  and  their  molecular  formulas  are,  for 
example,  O2,  H2,  C12,  N2.  (2)  The  molecular  weights  of  several 
metallic  elements  and  certain  gaseous  elements  are  identi- 
cal with  their  atomic  weights ;  that  is,  the  molecule  consists 
of  one  atom,  and  the  molecular  formula  is  the  same  as  the 
atomic  symbol,  e.g.  Na,  K,  Zn,  Hg,  Cd  (cadmium),  A  (argon), 
He  (helium),  and  Ne  (neon).  (3)  The  molecular  weights  of 
certain  elements  decrease  with  rise  of  temperature ;  e.g.  at 
lower  temperatures,  molecules  of  iodine,  sulphur,  and  phos- 
phorus are  represented  by  I2,  S8,  and  P4,  and  at  higher  tem- 
peratures by  I,  S2,  and  P2.  At  intermediate  temperatures 
the  two  kinds  of  molecules  are  in  equilibrium. 

Molecular  Equations.  —  Since  reactions  between  gases  are 
between  molecules,  equations  representing  such  reactions 
should  be  written  in  the  molecular  form.  For  example, 
since  a  molecule  of  hydrogen  has  the  formula  H2  and  of 
oxygen  the  formula  O2,  the  molecular  equation  for  the  forma- 
tion of  water  vapor  from  hydrogen  and  oxygen  is  — 

2  H2  +  02  =  2  H2O 

It  is  read  thus :  Two  molecules  of  hydrogen  unite  with  one 
molecule  of  oxygen  to  form  two  molecules  of  water  vapor. 
Since  most  elementary  gases  consist  of  molecules,  such 
equations  correctly  represent  the  actual  substances  in- 
volved. It  should  be  noted,  however,  that  the  proportions 
by  weight  are  the  same  as  in  the  simpler  or  atomic  form 
of  the  equation.  Molecular  equations  are  sometimes  called 
volume  equations  or  gas  equations.  Thus,  the  equation 

H2  +  C12  =  2  HC1 


252  INORGANIC  CHEMISTRY 

means  that  one  volume  each  of  hydrogen  and  chlorine  unite 
to  form  two  volumes  of  hydrogen  chloride,  or  more  fully 

H2         +         C12  2HC1 

2  gm.                        71  gm.  73  gm. 

22.4  1.                       22.4  1.  44.8  1. 

1  volume  1  volume  2  volumes 

1  molecule  1  molecule  2  molecules 

The  equation  may  be  written  in  the  latter  form  because 
equal  numbers  of  molecules  represent  equal  volumes.  It  is 
very  convenient  to  remember  that  in  molecular  or  gas  equa- 
tions one  molecule  represents  one  volume  of  a  gas. 

VALENCE 

A  classification  of  the  formulas  of  chemical  compounds 
shows  certain  regularities.  Instead  of  a  vast  number  of 
unrelated  formulas,  there  are  really  only  a  few  groups, 
especially  if  attention  is  confined  to  inorganic  compounds. 
Take,  for  instance,  some  binary  compounds  of  hydrogen. 
They  fall  into  the  following  classes :  — 

HC1  H2O  NH3  CH4 

HBr  H2S  PH3  SiH4 

HI  AsH3 

HF  SbH3 

Again,  if  some  oxides  are  considered,  we  have  the  follow- 
ing: — 

Na2O         CuO         A12O3         MnO2         P2O5         SO3         I205 
K2O          MgO         P2O3          SO2  As2O5       CrO3 

Ag2O         CaO         As2O3        CO2 

We  might  also  group  certain  acids  and  salts  thus :  — 
HC1  H2SO4  HNO3  H3PO4 

NaCl  Na2SO4  NaNO3  Na3PO4 

CaCl2  CaSO4  Ca(NO3)2  Ca3(PO4)2 

A1C18  A12(S04)3  Bi(N03)3  Mg3(P04), 


ATOMIC  AND   MOLECULAR  WEIGHTS  —  VALENCE      253 

Similarly,  there  are  the  following  groups  of  bases :  — 

NaOH  Ca(OH)2  A1(OH)8 

KOH  Ba(OH)2  Bi(OH)3 

If  these  lists  were  to  be  extended,  similar  regularities 
would  be  revealed.  A  careful  examination  of  these  formu- 
las leads  to  two  conclusions.  (1)  Atoms  of  elements  differ 
in  the  number  of  atoms  or  atomic  groups  of  the  other  ele- 
ments with  which  they  combine.  Thus,  in  the  first  list 
one  atom  each  of  chlorine,  bromine,  iodine,  and  fluorine 
combines  with  one  atom  of  hydrogen;  one  atom  of  oxygen 
and  of  sulphur  combines  with  two  of  hydrogen;  and  so  on 
through  all  the  lists.  (2)  An  atom  of  certain  elements 
unites  with  only  one  atom  or  atomic  group  of  certain  other 
elements,  with  only  two  atoms  or  two  atomic  groups  of  cer- 
tain others,  etc.  Thus,  one  atom  of  calcium  (Ca)  combines 
with  one  atom  of  oxygen,  with  two  of  chlorine,  with  one  SO4- 
group,  and  with  two  NOa-groups.  Hence  we  conclude  that 
atoms  of  the  elements  have  a  definite  and  characteristic  com- 
bining capacity.  The  number  which  expresses  the  maxi- 
mum combining  capacity  of  an  atom  of  an  element  is  called 
the  valence  of  the  element. 

Determination  of  Valence.  —  The  valence  of  an  element  is 
found  by  dividing  the  atomic  weight  by  the  equivalent 
weight.  It  will  be  recalled  that  elements  combine  in  the 
ratio  of  their  equivalent  weights.  That  is,  if  the  per  cent  of 
each  element  in  a  compound  is  restated  so  that  these  per 
cents  become  the  number  of  grams  which  unites  with  or 
replaces  eight  grams  of  oxygen,  the  composition  of  the  com- 
pound is  expressed  in  terms  of  the  equivalent  weights  of 
the  constituent  elements.  For  example,  magnesium  and 
oxygen  combine  in  the  ratio  of  60  and  40  per  cent  respec- 
tively (in  round  numbers).  If  for  the  40  per  cent  we  substi- 


254 


INORGANIC   CHEMISTRY 


tute  8,  then  the  60  per  cent  becomes  12,  i.e.  the  equivalents 
of  oxygen  and  magnesium  respectively.  We  have  seen, 
however,  that  atomic  weights,  not  equivalent  weights,  are 
the  weights  used  in  chemistry  for  expressing  composition. 
The  number,  therefore,  which  expresses  the  maximum  com- 
bining capacity  of  an  atom  of  an  element  is  merely  the 
factor  by  which  the  equivalent  weight  is  multiplied  to  con- 
vert it  into  the  atomic  weight.  To  find  the  valence,  it  is 
only  necessary  to  divide  the  atomic  weight  by  the  equivalent 
weight.  In  brief, 

Valence  =      Atomic  Weight 
Equivalent  Weight 

Tables   of   Valence.  —  Pursuing   this  method   of  finding 
valence,  we  obtain  the  following  table :  — 

A.  TABLE  OF  VALENCE  OF  COMMON  ELEMENTS 


ELEMENT 

SYMBOL 

VALENCE 

Aluminium    . 

Al 

III 

(Ammonium) 

(NH4) 

I 

Antimony 

Sb 

III  in  antimonous  and  V  in  antimonic  compounds 

Arsenic     .     . 

As 

III  in  arsenious  and  V  in  arsenic  compounds 

Barium      .     . 

Ba 

II 

Bismuth    .     . 

Bi 

III 

Boron  .     .     . 

B 

III 

Bromine    .     . 

Br 

I  in  hydrobromic  acid  (HBr)  and  bromides 

Cadmium  . 

Cd 

II 

Calcium     .    . 

Ca 

II 

Carbon      ,     . 

C 

IV 

Chlorine    .     .* 

Cl 

I  in  HC1  and  chlorides  ;  V  in  chlorates 

Chromium     • 

Cr 

III;  VI  in  chroinates  and  dichromates 

Cobalt  .     .    . 

Co 

II 

Copper      .     . 

Cu 

I  in  cuprous  and  II  in  cupric  compounds 

Fluorine    .     . 

F 

I 

Gold      .     .     . 

Au 

I  in  aurous  and  III  in  auric  compounds 

Hydrogen 

H 

I 

ATOMIC  AND   MOLECULAR  WEIGHTS  —  VALENCE      255 
A.   TABLE  OF  VALENCE  OF  COMMON  ELEMENTS  —  Continued 


ELEMENT 

SYMBOL 

VALENCE 

Iodine  .     .     . 

I 

I  in  hydriodic  acid  (HI)  and  iodides 

Iron      .     . 

Fe 

II  in  ferrous  and  III  in  ferric  compounds 

Lead     .     .     . 

Pb 

II  ;  IV  in  Pb02 

Lithium     .     . 

Li 

I 

Magnesium    . 

Mg 

II 

Manganese     . 

Mn 

II  ;  IV  in  MnO2  ;  VI  in  inanganates  ;  VII  in 

permanganates 

Mercury    .     . 

Hg 

I  in  mercurous  and  II  in  mercuric  compounds 

Nickel  .     .     . 

Ni 

II 

Nitrogen   .     . 

N 

1  in  N2O  ;  III  in  N2O3  and  nitrites  ;  V  in  N2O6, 

HNO3,  and  nitrates 

Oxygen     .     . 

0 

II 

Phosphorus   . 

P 

III  in  P203  ;  V  in  P2O6,  H3P04,  and  orthophos- 

phates 

Platinum  .     . 

Pt 

IV 

Potassium 

K 

I 

Silicon  .     .     . 

Si 

IV 

Silver    .     .    . 

Ag 

I 

Sodium      .     . 

Na 

I 

Strontium      i 

Sr 

II 

Sulphur    .     . 

S 

II  in  H2S  and  sulphides  ;  VI  in  S03,  H2SO4,  and 

sulphates 

Tin  .... 

Sn 

II  in  stannous  and  IV  in  stannic  compounds 

Zinc      .     .     . 

Zn 

II 

It  is  customary  to  call  valence  a  property  possessed  by 
atoms  of  elements.  But  groups  of  atoms  of  different  ele- 
ments often  act  chemically  like  individual  atoms  of  ele- 
ments (i.e.  they  pass  as  a  whole  from  compound  to  com- 
pound) ;  hence  it  is  customary  to  assign  a  valence  to  atomic 
groups  like  SO4,  NO3,  etc.  These  groups  are  often  called 
radicals,  and  are  analogous  to  NH4  and  OH,  Tabulating 
the  common  radicals  and  a  few  elements,  we  have :  — 


256 


INORGANIC  CHEMISTRY 


B.    TABLE  OF  VALENCE  OF  CERTAIN  ELEMENTS  AND  RADICALS 


CHKMICAL  NAME  OF  COMPOUND  WHICH  CONTAINS 
ELEMENT  OK  KADICAL 


SYMBOL  OR 

FollU  TLA 


VALENCE 


Acetate C2H3O2 

Bromide Br 

Carbonate  (normal) CO3 

Carbonate  (acid) HCO3 

Chlorate C103 

Chloride Cl 

Chromate Cr04 

Cyanide CN 

bichromate Cr2()7 

Ferricyanide Fe(CN)6 

Ferrocyanide Fe(CN)6 

Fluoride F 

Hydroxide OH 

Iodide I 

Manganate MnO4 

Nitrate NO3 

Nitrite NO2 

Oxide O 

Permanganate Mn04 

Phosphate  (ortho) PO4 

Silicate  (meta) SiO3 

Sulphate SO4 

Sulphate  (acid) HSO4 

Sulphide S 

Sulphite  (normal) S03 

Sulphite  (acid) HS03 


I 

I 

II 

I 

I 

I 

II 

I 

II 

III 

IV 


I 


II 

I 

III 

II 

II 

I 

II 

II 

I 


It  is  obvious  from  these  tables  that  the  valence  of  some 
elements  varies.  Many  elements  have  a  fixed  valence,  but 
several  of  the  common  elements  have  a  valence  which  is 
determined  by  the  element  with  which  they  are  united  and 
the  conditions  under  which  the  union  occurred.  The 
valence  of  an  element  is  never  less  than  1  nor  more  than  8. 
The  valence  of  most  radicals  is  fixed. 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE    257 

Valence  Terms.  —  Elements  and  groups  with  the  valence 
I  are  called  monads  or  univalent  elements,  and  those  with 
the  valence  II  are  called  dyads  or  bivalent  elements ;  simi- 
larly, elements  or  groups  which  have  the  valence  III,  IV,  V, 
VI,  and  VII  are  called  respectively  triads,  tetrads,  pentads, 
hexads,  and  heptads,  and  the  corresponding  terms  are  triva- 
lent,  quadrivalent,  quinquivalent,  hexavalent,  and  heptav- 
alent. 

Valence  applied  to  Combination  and  Displacement.  — 
(1)  Elements  which  have  the  same  valence  combine  with  each 
other  atom  for  atom.  Thus,  one  atom  of  sodium  combines 
with  one  atom  of  chlorine,  and  one  atom  of  magnesium  com- 
bines with  one  atom  of  oxygen.  (2)  On  the  other  hand,  ele- 
ments having  a  different  valence  usually  combine  with  each 
other  so  that  the  total  valence  of  the  atoms  of  each  element 
in  a  molecule  is  equal,  i.e.  in  many  compounds  the  total 
combining  capacity  of  each  element  is  equal.  This  condi- 
tion is  often  described  by  calling  the  compound  saturated, 
or  as  one  having  a  balanced  valence.  Thus,  two  atoms  of 
hydrogen  combine  with  one  of  oxygen  to  form  one  molecule 
of  water  (H20) ;  two  atoms  of  aluminium  each  having  the 
valence  III  combine  with  three  atoms  of  oxygen  each  hav- 
ing the  valence  II  to  form  one  molecule  of  aluminium  oxide 
(Al2Os) ;  one  atom  of  carbon  having  the  valence  IV  com- 
bines with  two  atoms  of  sulphur  each  having  the  valence  II 
to  form  one  molecule  of  carbon  disulphide  (CS2).  (3)  The 
above  rules  apply  to  atomic  groups.  Thus,  one  NH4-group 
combines  with  one  OH-group  to  form  one  molecule  of  ammo- 
nium hydroxide  (NH4OH) ;  two  NH4-groups  combine  with 
one  SO4-group  having  the  valence  II  to  form  one  molecule 
of  ammonium  sulphate  ((NH4)2S04) ;  and  three  atoms  of 
calcium  each  having  the  valence  II  combine  with  two  PO4- 
groups  each  having  the  valence  III  to  form  one  molecule  of 
calcium  phosphate  (Ca3(PO4)2). 


258  INORGANIC   CHEMISTRY 

It  is  convenient  to  interpret  valence  not  only  from  the 
standpoint  of  combination,  but  also  from  the  standpoint  of 
displacement.  That  is,  just  as  atoms  and  groups  of  the 
same  valence  combine  unit  for  unit,  or  those  of  different 
valence  combine  so  that  the  valences  of  each  element  or 
atomic  group  balance,  so  also  atoms  and  atomic  groups  dis- 
place each  other  —  unit  for  unit  if  the  valence  is  the  same, 
or  equivalently  if  the  valence  is  different.  For  example, 
when  zinc  displaces  hydrogen  in  hydrochloric  acid,  one  atom 
of  zinc  having  the  valence  II  displaces  two  atoms  of  hydro- 
gen each  having  the  valence  I,  and  the  formula  of  the  result- 
ing zinc  chloride  is  ZnCl2 ;  similarly,  one  atom  of  sodium  having 
the  valence  I  displaces  one  atom  of  hydrogen  from  water,  and 
the  formula  of  the  resulting  sodium  hydroxide  is  NaOH. 

Valence  and  Formulas.  —  The  chief  value  of  valence  in  a 
course  in  inorganic  chemistry  is  the  assistance  it  gives  in 
writing  formulas.  If  it  were  necessary  to  remember  the 
formula  of  every  compound,  the  study  of  chemistry  would 
be  tedious  and  almost  hopeless.  The  application  of  valence 
to  formula  writing  can  be  best  illustrated  by  several  ex- 
amples. Suppose  we  wish  to  write  the  formula  of  mag- 
nesium chloride.  From  Table  A  the  valence  of  magnesium 
is  found  to  be  II,  and  from  Table  B  the  valence  of  chlorine 
in  chlorides  is  found  to  be  I ;  therefore,  it  is  necessary  to  take 

two  atoms  of  chlorine  for  one  of  magnesium  in  order  to 

ii  ii 

balance  the  valence,  and  the  formula  is  MgCl2.  It  is  help- 
ful to  write  the  valence  as  shown  in  the  formula  just  given, 
though  such  a  device  may  be  abandoned  after  the  valence 
of  each  element  and  group  has  been  learned.  Again,  sup- 
pose we  wish  to  write  the  formula  of  lead  nitrate.  As  in 
the  previous  case,  from  Table  A  the  valence  of  lead  is  found 
to  be  II,  and  from  Table  B  the  valence  of  the  NO3-group  in 
nitrates  is  found  to  be  I  ;  therefore  it  is  necessary  to  take 


ATOMIC  AND  MOLECULAR  WEIGHTS  —  VALENCE     259 

two  NO3-groups  to  balance  the  valence  of  the  lead,  and  the 

formula   is   Pb(NO3)2.     Similarly,    the   formula   of   sodium 

ii     ii 

sulphate  is  Na2SO4  because  two  atoms  of  sodium  each  hav- 
ing the  valence  I  are  needed  to  balance  the  valence  II  of 

the  SO4-group  ;  and  the  formula  of  aluminium  hydroxide  is 

in    in 

A1(OH)|  because  three  OH-groups  each  having  the  valence 

I  are  needed  to  balance  the  valence  of  one  atom  of  aluminium 
having  the  valence  III.  Furthermore,  suppose  we  wish  to 
write  the  formula  of  a  compound  formed  by  the  interaction 
of  aluminium  and  hydrochloric  acid.  The  valence  of  alumin- 
ium is  found  from  Table  A  to  be  III,  and  that  of  hydrogen 

to  be  I.     Therefore  one  atom  of  aluminium  displaces  three 

in  in 
atoms  of  hydrogen,  and  the  formula  is  A1C13.     Finally,  the 

formulas  of  acids,  although  usually  learned  by  continued 
observation,  can  readily  be  written  by  utilizing  Table  B. 
Acids  may  be  regarded  as  compounds  formed  by  the 
union  of  hydrogen  with  radicals  and  elements  in  Table  B. 
That  is,  to  write  the  formula  of  an  acid  it  is  only  necessary 
to  prefix  the  proper  number  of  hydrogen  atoms  to  the  atom 
or  group  in  the  table,  remembering,  of  course,  the  rules  for 

naming  acids.     Thus,  the  formula  of  chloric  acid  is  HC1O3, 

because  the  valence  of  C1O3  is  I,  and  this  valence  is  balanced 

i    i 
by  one  atom  of  hydrogen;   similarly,  nitrous  acid  is  HNO2, 

in  in  ii    n 

phosphoric  acid  (ortho)  is  H3PO4,  silicic  acid  is  HgSiOs. 

The  student  is  advised  to  write  the  formulas  of  many  com- 
pounds until  valence  and  its  application  become  thoroughly 
familiar.  (See  Exercises  at  the  end  of  this  chapter.) 

Representation   of  Valence.  —  One   way  of  representing 

the  valence  of  elements  and  radicals  has  been  given,  viz. 

ii    i      ii 

O,  Cl,  SO4.     Sometimes  short  lines  are  used,  e.g.  H  — ,  O  = ,  or 


()         ,    AI          ,    He          II     1  1  IK'S    .MIC   II  cd    !(>   in.  li.-.-i  I  <•   Y.'lltMHT  ill 

eompound   .   .1     in."!.'  hue  ;in   \\rrs  for  1  wo  rlminils.      Thus. 

III.     I  ......  ill.  i    III'   \\.Ml.T   ill        O        ||    ,.,i  1,,-j     !  li.in    ||  ()  ||. 

/N 

Siiinl.i  1  1  \  ,  .immonia  jra:;  is    N      II   MIX!  eah'ium   hydroxide  is 

\H 

/OH 

CM  .      The    I,  -in-Ill.    ,.f    ill.-    lines    li;iv.-    no   sitfniliranrr, 


(he  lines  do   iml    ivpi  vsrnl    mien   n\    of  at  t  I'.'icl  ion   l>«- 
I\\««-M    ili«-   iitoms,      Such    formula.-.   ;u.-   r:ill«'d    .-.(  i  m  (  u  i  .d   or 
Kuiphir    totmuhis,    :ind    HI.-N     ol'h-n    -..-i\.-    !,.    sho\\     \\li.Ml    is 
known  .il.niii    ili«-  .11  iMiaMMiitMi!   of  the  nloins  in  a  inoleruhv 


Apparent  Exceptions.  -    The  formulas  of  some  r<.m|>oimds 
d..  no  i  conform  Id  (he  ronee|i{ion  ..f  yaleure  diseusstul  ;il>oye, 

t,^,  carbon  monoxide  (C(>K  .•airimn  carbide  iCat1,,),  mntf- 

nefic    .....i    oxide    (  I<V,<  >    ,    nei^lylene    ((  '  \\^.    and    elln  leno 
((  ',11,).      Tlieii   inlerprelalion  musl   h«-  soii^li!  elsewhere. 

I'uom  I-MM    \M.    I\I.-K,  -ISI.-M 

I     Til.-  \.i|...i    .l,-itsitit\s  «»f   tM'fiaiti  KUHOH  (rt«forn-.l  i.«  ..-.\j-.-iO  MI-«« 
M  .  l..||,,u  .      (,.»  liyilr.irlilorii-  :«nd   I    II  .   iM  ,  l.|,.nn,-  "  '.'  I  S.   (.  >  .n.un,  M.I.I 

.Aaiafl,  («/)  nttrtiKt^n  .NTrt,  (e)  iteam  .SOW.     t'uloulal^  thr  modular 
\vt»l«ht  t»f  oiioh. 

'J.    (^tUtMlltito    tlif    -.linplr    I     I.Miiiul.i    .'I     (lie    t'oiupoimils    \\hi.h   li.ix.- 

Uw  Inditwtt'tl  cninpnNititui  II      17.  01  V.   (/-to     ;;o. 

Ko  (Ihtn)  «  70;    (r)  It  ~=  1,  0  *  I',  is    (poU    turn)  *  39,  O  »  48. 
\  liltM1  of  milplmi.Mi    "\i.l.   .-..i     (HO   I  Weigh     '  M-.  '  .-.ni       Wluit. 
IH  (ht»  iuul«MMilur  woighl  uf  (his  tH>in|uMin.l  ; 

I     If   lf»00  oo.  t»f  ourbiMi   in.  -ii.  -\jilo  gns   (00)   Weifh    I  SMIO  gnu, 
\\li:ii  i.  tlu-  m.iltMMiltir  \voi)thi  «•!  il>.-  tMuupuuiuil 

ft,  t^vJoulaU*  (!»«>  in.'!.-,  ul.  u   i,  M  nuil.  i  .'i  (h,-  compounds  correspond 
(UK  to  ill.-  i,  -II-  '\viu«  dutti:    («)  0  »  78,8,  H  •  8,7,  N  ss  17,1,  vapoi 
density  «  fl.Oa  ;   (h)  0  •  02,3,  H  »  7.7,  vwp."   dentito        '  i 
C  »  30,6,  H  «  0,7,  ()  <•  f»a,  I,  vA|K>r  den»ity  «  l.DOi.      |  \  ..    <  density 
ni 


ATOMIC    AND    MOU'XTLAU    \\TICIITS        YAI.MMT     -.Mil 
ii     \\  II.H   \..luni,     of  i.i,  I,H  .  .iii.l  products  are  represented  by  tho 

.•qu.iiion .  („)  il  i  ci.  '2  IH  M.  (/>)  •.'  M.  i  o.  •.•  i i.o.  (r)  ;;  11, 
I  N-;  '2  N I li,  (<n  Nj  I-  Oi— 2  No.  <« )  •.'  NO  |-  Oi  =  2  No." 

7.  II'  «H.1(»  >vm.  of  melalli.  il\,i  \i,l.l  I'l  l'>'».{  kMM.  of  puro 
silver  chloride,  calculate  I  ho  atomic  wen-lit  of  chlorine 

S.  Suppose  -I.S(>1  1 1  Kin.  of  ferric  oxide  il-V.O,)  \  icld  •  :i',MM>.">  »riu. 
of  iron.  \\li:il  is  I  lu«  ;i  1 01111,  \\iMv.lil  ol'iion' 

«.).  A  .nun'  lli.'it  I  ,().>!.:;  rin  •'!'  inmliiim  hromido  (SrHi'.j)  ro- 
(jiiiro  I.IS/OV  fjii-  «»•'  sil\«'r  lo  prooipilMto  Iho  hromiiu*  as  A;-.Ui. 
NNIuil  is  Ilio  Mlonih-  \\civ.lil  of  ;!i-t)ii!iiiin  ? 

10.  If  l'.»  .>,  I  '(»  .'iii.  ol'  iiHMvuri.'  «-lilon.l««  (llkr<M9)  yioM   I  I.-KMKW 
IMII    «.l   in, -i,  HI  \.  \\luii   is  I  ho  jiloinic  \V(<i>';lit  <>!'  MUM-CIIPN  ? 

11.  A  litor  of  llu«  inoiinloiiiic  K-'»S  holiiim   \V«MK|IS  .17SH  ^m     t:l1 

0  '  ( '.  iind  7(>()  mm.).      NVIi.'H  r    I  h«-  .Homic  wiM^lil.  of  liolium  ? 

rj.    (,i)   .sr.oc*  i-.m.  of  |I\.|I-.»IM«H  , oiuKini'ii  \\iih  c>s  •.';>();{;{  ^\\\. 

..I'  1. 1. .IIIIIH-  (  '.il.  ill. il.-  III.-  .H.. mi. •  \\.-i--hl  ..I'  1. 1. .linn.-  (/»)  I  .M.M 
I'.in.  of  (-:ilriiim  hroniido  (('.-iHi-.)  i-»M|inr.'  l.'.K"»0'jr>  £\\\.  of  >.tl\.-r  lo 
pnM'ipilnl(>  Ilio  hromino  MS  A>'.Hr.  \\'|IM(  is  (ho  iitomio  woi^lil  «>f 

r.-ll.'Ullii  " 

I  :  i.f)  II'  I...'.*..!)/  -MM.  of  phosphorus  I  rirhloriilo  iIMM;,)  >;iv«> 
1-l.liSllS  K'H-  of  silv»»r  chlorido.  wluil  is  Iho  atomic  wiM^lil.  of  phon- 
phoniH?  (l>)  Tin1  moItuMilar  wtM^lil  of  curium  disulphidc  is  7(>,  tuul 
I(K)  p. H  i  yi(>ld  NI  'I  parts  of  sulphur;  (.ho  spocilic  (mat.  of  sulphur 
in  .17S.  \Yh:il  is  lh.'  :ili>nu<'  \\ciKhl.  of  sulphur,  ami  how  mnuy 
atoms  of  siilpluii  in  .1  molecule  of  tho  (*arl)on  disulphid.  ' 

I  I  ir-.M'JSI  KIM.  of  N-jO^-ix".!;  ISOKm.of  Nu.  what  is  the  atomic 
u  .  i"  li  I  ol  nil  ri>K<*ll  ? 

I.'      A  .TI-IMIII   \vciivhl   <>!'  copper  oxide,  when  henled  it)  a  current 

01  IIN.IIO.-.MI.  lost  .r)U.7S«J  Kin.  ol   O\\K«MI  :ind  1'itrmod  (17.2S2  KHI-  <>f 
\\.-iier       t.i)    If   O       Hi.    \\hMl    is   Iho  Mhimir    xxei^lil.  of    hydrogen? 
(M   If  II        1.  \\li.Ml    is  (he  Mloinir  weJKlit.  of  ox.vgiMi  ? 

\(\.  (a)  \\  h.-il  is  lhea.tomi««  weiKlH  of  mercury  if  the  specific  heat, 
is  0.():;-J':'  (M  Of  le.-id  if  the  specific  IICM!  is  ().();{ I? 

17.    («i)  To  foi*m  a  certniu  compound,   I  K'"-  «>f  carbon  mid  '.'  <>i'(> 
mi    ol  o\\;-.eii  :ire  needed        I  Is  \  jipor  diMisil.y  n»forre«l  (<»  ONVKCII  i. 
I    ••   •       \\IIM!    is  its  molecular  formula  ?      (!>}    A    compound  ol    car 
bon  and   hydrogen  contains   1-1.21)  per  cent    of   hydroKen,  and    I    I. 
weii'lis    I  LT)  Kin.      ( 'alculate  il  ;  molecular  formula. 

IS.  If  1.2  tfin.  of  a  substance  dissolved  in  21. ,r>  Km-  °f  ^aler 
lower  the  friM»/,inK  |)oinl  !.().»,  \\|i;l|  (s  the  inoloiMiltir  weight-  of  the 
Hnbstan,  , •" 


262  INORGANIC  CHEMISTRY 

19.  Write  the  formula  of  the  chloride  of  potassium,  sodium,  sil- 
ver, copper  (ous),  copper  (ic),  mercury  (ous),  mercury  (ic),  iron 
(ous),  iron  (ic),  cadmium,  zinc,  tin  (ous),  tin  (ic),  calcium,  barium, 
magnesium,  strontium,  cobalt,  nickel,  bismuth,  aluminium,  carbon, 
ammonium,  antimony  (two),  phosphorus  (two).     (Two  means  -ous 
and  -ic.) 

20.  Write  the  formula  of  the  sulphide  of  K,  Na,  Ag,  Cd,  Zn,  Ca, 
Ba,  Co,  Ni,  NH4,  As  (two),  Sb  (two),  Sn  (two),  Pb,  Fe  (ous). 

21.  Write  the  formula  of  the  sulphate  of  K,  Na,  Ag,  Cu,  Fe  (ous), 
Cd,  Zn,  Pb,  Ca,  Ba,  Co,  Mg,  Ni,  Mn,  Sr,  Cr,  Al,  NH4. 

22.  Write  the  formula  of  the  nitrate  as  in  Problem  21. 

23.  Write  the  formula  of  the  hydroxide  of  NH4,  Al,  K,  Na,  Ca, 
Mg,  Ba,  Fe  (ic),  Ni,  Co,  Cd,  Zn,  Cr,  Bi. 

24.  Write  the  formula  of  the  potassium  compound  of  the  ele- 
ments and  radicals  given  in  Table  B  (under  Valence,  above).     Do 
the  same  with  Ca  and  Mg. 

25.  Write  the  formula  of  copper  acetate,  lithium  chlorate,  man- 
ganese   dioxide,    ammonium    fluoride,    sodium    silicate,  potassium 
manganate,  barium  phosphate,  zinc  iodide,  ammonium  chromate, 
silver  chromate,  lead  acetate,  cobalt  nitrite,  magnesium  oxide. 

26.  Write  the  formula  of  aluminium  oxide,  auric  chloride,  fer- 
rous bromide,  ferrous  carbonate,  mercurous  nitrate,  platinum  chlo- 
ride, aluminium  phosphate,  calcium  fluoride,  potassium  cyanide. 

27.  Write  the  formula  of  phosphoric  acid,  silicic  acid,  sulphur- 
ous acid,  nitrous  acid,  chromic  acid,  acetic  acid,  hydriodic  acid. 

28.  By  the  Victor  Meyer  method  .1561  gm.  of  a  compound 
expelled  32.1  cc.  of  dry  air  at  20°  C.  and  744  mm.     Calculate  its 
vapor  density  and  its  molecular  weight. 

29.  What  volume  of  the  component  gases  can  be  obtained  by 
the  decomposition  of  6  1.  of  ammonia  gas? 

30.  What  volume  of  oxygen  is  used  up  when  20  cc.  of  acetylene 
burn  in  air?     (Equation  is  2  C2H2  +  5  O2  =4  CO2  +  2  H2O.) 

31.  (a)  One  volume  of  phosphorus  vapor  and  six  volumes  of 
chlorine  gas  form  four  volumes  of  phosphorus  trichloride  vapor; 
write  the  molecular  equation  for  this  chemical  change.     (6)  How 
many  liters  of  chlorine  would  be  needed  for  248  gm.  of  phosphorus  ? 

32.  Calculate  the  following  equivalent  weights:   (a)  .464  gm.  of 
zinc  gives  174  cc.  of  hydrogen  at  18°  C.  and  763  mm.     (6)  .075  gm. 
of  magnesium  gives  75.6  cc.  of  hydrogen  at  17°  C.  and  760.4  mm. 
(c)  .17  gm.  of  aluminium  gives  223.7  gm.  of  hydrogen  at  16°  C.  and 
754  mm. 


CHAPTER  XV 
Carbon  and  its  Oxides  —  Carbides 

Occurrence  of  Carbon.  —  Uncombined  carbon  is  found 
pure  in  nature  as  diamond  and  graphite  ;  in  a  more  or  less 
impure  state  it  occurs  as  coal  and  similar  substances,  which 
are  included  in  the  .term  amorphous  carbon.  Combined  with 
hydrogen  and  oxygen,  and  occasionally  with  nitrogen  also, 
it  is  an  essential  constituent  of  animal  and  vegetable  matter. 
Meat,  starch,  fat,  sugar,  wood,  cotton,  wool,  wax,  flour, 
albumen,  and  bone  are  familiar  examples  of  the  vast  num- 
ber of  natural  substances  which  contain  carbon.  It  is  also 
a  constituent  of  carbon  dioxide  and  of  the  carbonates,  such 
as  limestone,  chalk,  marble,  and  dolomite.  Illuminating 
gases,  gasolene,  kerosene  (and  other  products  obtained 
from  petroleum),  turpentine,  alcohol,  chloroform,  ether,  and 
similar  liquids  are  compounds  of  carbon.  Carbon  is  also  a 
constituent  of  thousands  of  manufactured  compounds,  such 
as  dyestuffs,  medicines,  and  perfumes. 

Diamond  is  pure  crystallized  carbon.  It  is  found  in  only 
a  few  places  in  the  earth.  When  taken  from  the  mine, 
diamonds  are  usually  rough-looking  stones;  some  are 
crystals,  some  are  rounded  like  peas,  and  many  are  irregu- 
lar; they  are  ground  into  special  shapes  and  polished  to 
bring  out  the  luster  and  make  them  sparkle  (Fig.  28).  The 
most  expensive  diamonds  are  colorless  and  without  a  flaw, 
and  are  said  to  be  "  of  first  water  " ;  diamonds  having  a 
yellow  tint  are  common,  and  occasionally  a  blue,  pink,  or 
green  one  is  found  ;  a  very  impure  variety  is  black. 

263 


264 


INORGANIC  CHEMISTRY 


Diamond  resists  the  action  of  most  chemical  reagents. 
It  has  the  high  specific  gravity  of  3.5,  and  is  one  of  the  hard- 
est substances.  It  is  brittle,  and  may  be  shattered  by  a  blow 
with  a  hammer. 


Crystal 


Rough 
FIG.  28.  —  Diamonds. 


Cut 


Diamonds  have  always  been  prized  as  gems  on  account  of  their 
beauty,  rarity,  and  permanence.  Besides  being  worn  as  jewels,  they 
are  used  to  cut  glass,  and  the  powder  and  splinters  (known  as  bort) 
are  used  to  grind  and  polish  diamonds  and  other  hard  gems.  The 
impure  variety  is  called  carbonado,  and  is  set  into  the  end  of  the 
"  diamond  drill,"  which  is  used  extensively  for  boring  artesian  wells 
and  drilling  hard  rocks. 

Diamonds  were  formerly  found  in  gravel  deposits  in  India,  and 
in  later  years  in  Brazil.  Since  1867,  however,  about  95  per  cent  of 
the  diamonds  of  commerce  have  come  from  South  Africa.  They 
occur  in  a  bluish  volcanic  rock  along  the  Vaal  River,  and  especially 
near  Kimberley.  Over  ten  tons  of  diamonds  have  been  found  in 
South  Africa  in  the  last  twenty-five  years  ! 

The  successive  investigations  of  Lavoisier,  Dumas,  and  Davy,  ex- 
tending from  1772  to  1814,  showed  that  diamond  is  carbon,  for  when 

pure  diamond  was  burned  in  oxy- 
gen, the  only  produce  was  carbon 
dioxide.  This  result,  which  ad- 
mits of  no  doubt,  has  been  verified 
by  many  famous  investigators. 
Diamonds  have  been  made.  In 
1893  Moissan  dissolved  pure  char- 
coal in  melted  iron,  and  cooled  the  molten  mass  in  water.  The 
surface  was  so  suddenly  cooled  that  a  tremendous  pressure  was 
exerted  by  the  expanding  iron  inside  the  crust.  This  pressure 
caused  the  cooling  carbon  to  crystallize  into  diamond.  The  crystals 


CARBON  AND  ITS  OXIDES  —  CARBIDES         265 

were  very  small,  most  of  them  were  black,  a  few  were  white,  but  all 
had  the  properties  of  the  diamond  (Fig.  29). 

Large  diamonds  have  a  fascinating  history,  since  most  of  them 
passed  through  many  hands  before  becoming  royal  jewels.  Until 
recently  the  largest  royal  cut  diamond  was  the  Orloff,  which  weighs 
194|  carats,  and  is  in  the  scepter  of  the  Czar  of  Russia.1  The  Kohi- 
noor,  which  now  weighs  about  106  carats,  is  one  of  the  crown  jewels 
of  England.  The  largest  diamond  ever  found  was  called  the  Culli- 
nan.  It  was  found  in  South  Africa  in  1905,  and  weighed  about 
3025  carats  (1.37  Ib.  avoir.).  Stones  cut  from  it  are  among  the 
English  crown  jewels. 

Graphite  is  a  soft,  black,  shiny  solid,  which  is  smooth 
and  greasy  to  the  touch.  Pure  graphite  is  entirely  carbon. 
It  occurs  native  in  large  quantities  and  in  many  places. 
One  variety  is  found  in  abundance  at  Ticonderoga,  New 
York.  Other  localities  are  Mexico,  Ceylon,  Siberia,  Ger- 
many, Austria,  and  Italy.  Sometimes  crystals  and  grains 
are  found,  but  it  usually  occurs  in  flaky  masses  or  slabs. 
Unlike  diamond,  graphite  is  a  good  conductor  of  electricity 
and  is  often  used  to  coat  molds  in  electrotyping.  It  is 
so  soft  that  it  blackens  the  fingers  and  leaves  a  black  mark 
on  paper  when  drawn  across  it.  This  property  is  indi- 
cated by  the  name  graphite,  which  is  derived  from  a  Greek 
word  (graphein)  meaning  to  write.  It  resembles  diamond 
in  its  insolubility  in  liquids  at  the  ordinary  temperature. 
Its  specific  gravity  is  about  2.2,  being  considerably  lighter 
than  diamond.  When  heated  intensely  in  a  current  of 
oxygen,  it  is  converted  into  carbon  dioxide  ;  but  it  can  be 
heated  to  a  very  high  temperature  in  the  air  without  under- 
going chemical  change.  Graphite  was  once  supposed  to 
contain  lead,  and  is  even  now  often  incorrectly  called  "  black 
lead"  and  plumbago.  It  is  used  to  make  stove  polish  and 
protective  paints,  as  a  lubricant  where  oil  would  be  de- 
composed by  the  heat,  as  the  principal  ingredient  of  the 

1  The  international  carat  weighs  200  milligrams. 


266  INORGANIC  CHEMISTRY 

graphite  crucibles  which  withstand  extremely  high  tempera- 
tures, and  in  making  electrodes  for  electric  furnaces  and 
electrolytic  apparatus.  Large  quantities  of  graphite  are 
consumed  in  the  manufacture  of  lead  pencils.  The  graphite 
is  ground  to  a  fine  powder,  mixed  with  more  or  less  clay, 
and  then  passed  through  perforated  plates,  from  which  the 
"  lead  "  issues  in  tiny  rods.  These  are  dried,  cut  into  proper 
lengths,  baked,  and  then  inserted  in  the  wooden  case. 

Molten  iron  and  other  metals  dissolve  carbon,  and  when 
the  metals  cool,  the  carbon  crystallizes  as  graphite. 

Large  quantities  of  graphite  are  now  manufactured  by  heat- 
ing a  special  grade  of  coal  or  of  coke  in  an  electric  furnace. 
(Fig.  20).  The  process  is  electrothermal,  and  yields  a  product 
that  is  exceptionally  suitable  for  electrodes.  Articles  of 
almost  any  size  and  shape  can  be  made  of  graphite ;  more- 
over the  graphite  is  very  compact  and  can  be  further  shaped 
by  tools. 

Amorphous  Carbon  is  a  broad  term,  including  all  varieties 
of  coal,  coke,  charcoal,  lampblack,  and  gas  carbon.  They 
are  the  non-crystalline  varieties  of  impure  carbon,  and  differ 
mainly  in  purity,  degree  of  fineness,  and  hardness.  The 
word  amorphous  means  literally  "without  form,"  and  is 
often  used  to  designate  soft,  powdery  substances,  but  more 
especially  those  which  are  uncrystallized. 

Coal  is  the  term  applied  to  several  varieties  of  impure 
carbon.  It  may  be  regarded  as  the  final  product  derived 
from  vegetable  matter  which  was  subjected  to  heat  and 
pressure  through  long  geological  periods.  Ages  ago  the 
vegetation  was  exceedingly  dense  and  luxuriant  upon  land 
slightly  raised  above  the  sea.  In  the  process  of  time  this 
vegetation  decayed,  accumulated,  and  slowly  became  covered 
with  sand,  mud,  and  water.  Owing  to  the  heat  of  the  earth 
and  the  enormous  pressure  of  the  overlying  deposits,  the 


CARBON  AND   ITS  OXIDES  —  CARBIDES          267 

complex  vegetable  matter  decomposed  slowly   into    carbon 
and  compounds  of  hydrogen  and  carbon  called  hydrocar- 


FIG.  30.  —  Section  of  part  of  the  earth's  crust  near  Mauch  Chunk, 
Pennsylvania,  showing  layers  of  coal. 

bons.  The  geological  and  chemical  changes  were  repeated, 
and  as  a  result  we  find  in  the  earth  layers  or  seams  of  car- 
bonaceous matter  varying  in  thickness  and  composition 
(Fig.  30).  These  are  the  coal  beds.  Coal  beds  contain  proofs 
of  their  vegetable  origin,  viz.  impressions  of  vines,  stems, 
leaves  of  plants,  and  similar  vegetable  substances  (Fig.  31). 


FIG.  31.  —  Fossil  found  in         FIG.  32.  —  Section  of  coal  as  seen  through 
a  coal  bed.  a  microscope. 

A  thin  section  of  coal  examined  through  a  microscope 
reveals  a  distinct  cellular  structure  characteristic  of  vege- 
table matter  (Fig.  32).  The  transformation  of  vegetable 
matter  into  coal  was  an  exceedingly  slow  process,  espe- 


268  INORGANIC   CHEMISTRY 

cially  where  the  rocks  were  undisturbed.  In  some  locali- 
ties the  process  was  hastened  by  deep-seated  changes 
incidental  to  the  upheaval  of  the  rocks  during  mountain 
building ;  the  gaseous  products  escaped  almost  com- 
pletely and  left  the  carbon,  together  with  a  small  pro- 
portion of  mineral  matter.  Consequently  we  find  several 
varieties  of  coal  in  the  earth's  crust.  There  are  three  main 
classes  —  anthracite,  bituminous,  and  lignite,  though  com- 
mercially several  subclasses  are  variously  designated.  They 

LIGNITE,  CROCKETT,  TEX. 

Fixed  Carbon  Moisture  Ash  Volatile  Matter 

LIGNITE,   GLENDIVE,  MONT. 

F.C.  »T  A.  V.M. 

SUB-BITUMINOUS,   LAFAYETTE.  COLO. 

F.C.  M.  A.  V.  M. 

SUB-BITUMINOUS,  GALLUP.  N.MEX. 

F.C.  M.  A.  V.M. 

BITUMINOUS,  CARTERVILLE.  ILL. 

F.C.  M.      A.  V.M. 

SEMI-BITUMINOUS,   POCAHONTAS,  W.  VA. 

F.C.  M.          A.  V.M. 
SEMI-ANTHRACITE,   SPADRA,  ARK. 

F.C.  M.        A.         V.M. 

ANTHRACITE,  PENNSYLVANIA 

F.C.  M.  A. V.M 

FIG.  33.  —  Diagram  showing  the  progressive  change  in  composition  of 
coal.     (From  National  Geographic  Magazine.) 

differ  in  composition.  That  is,  they  contain  different  pro- 
portions of  moisture,  ash  or  mineral  matter,  volatile  matter 
or  compounds  of  carbon  and  hydrogen  driven  off  by  careful 
heating,  and  fixed  carbon  or  carbon  left  after  the  removal 
of  the  volatile  matter  and  ash.  The  progressive  change  in 
composition  from  a  very  poor  lignite  to  the  best  anthracite 
is  shown  in  Figure  33.  The  diagram  also  shows  the  com- 


CARBON  AND   ITS  OXIDES  —  CARBIDES          269 

position  as  determined  by  proximate  analysis,  i.e.  a  chemical 
analysis  designed  to  find  the  per  cent  of  the  components 
mentioned  above  and  not  the  actual  per  cent  of  carbon, 
hydrogen,  etc.  From  the  diagram  it  is  seen  that  the  fixed 
carbon  (F.C.)  varies  from  about  20  per  cent  in  the  lignite 
to  about  90  per  cent  in  the  anthracite,  while  the  volatile 
matter  (V.M.)  varies  inversely  as  the  fixed  carbon,  being 
greatest  in  the  lignite  and  least  in  the  anthracite.  The  per 
cent  of  moisture  (M.)  diminishes  from  the  lignite  to  the 
anthracite.  The  per  cent  of  ash  (A.)  is  variable. 

The  properties  and  uses  of  coal  differ  with  the  composi- 
tion. Anthracite-  coal  is  hard  and  lustrous.  It  ignites 
with  difficulty,  burns  with  a  slight  or  no  flame,  and  pro- 
duces an  intense  heat.  It  is  used  mainly  for  domestic  pur- 
poses, —  heating  and  cooking,  —  especially  in  the  eastern 
United  States.  Bituminous  or  soft  coal  burns  with  a  smoky 
flame,  and  in  burning  produces  considerable  volatile  matter  ; 
some  varieties  form  coke  when  heated.  It  is  used  to  make 
illuminating  gas,  coke,  and  as  a  fuel  for  generating  steam. 
Lignite  or  brown  coal  is  the  least  valuable  as  fuel.  The  use 
of  coal  as  fuel  depends,  of  course,  on  the  fact  that  consider- 
able heat  is  liberated  when  it  burns;  that  is,  it  has  calorific 
value.  The  following  table  shows  the  calorific  value  of 
well-known  commercial  grades  of  the  three  classes  of  coal 
in  calories  per  gram,  i.e.  the  number  of  calories  liberated 
when  1  gm.  is  burned  freely:  — 

CALORIFIC  VALUE  OF  COAL 


CLASS 

CALORIES  PER  GRAM 

Anthracite      

7724 

Bituminous    

8768 

Lignite  

4530 

270 


INORGANIC   CHEMISTRY 


Coal  is  widely  distributed  in  the  crust  of  the  earth,  but  the  deposits 
vary  in  extent  and  quality.  It  underlies  about  one  sixth  of  the  area 
of  the  United  States,  the  anthracite  variety  covering  less  than  five 
hundred  square  miles  in  eastern  Pennsylvania  (Fig.  34).  The  United 


FIG.  34.  —  Map  showing  coal  areas  in  the  United  States.  The  black 
areas  are  anthracite  and  bituminous  ;  the  shaded  areas  are  lignite.  (From 
National  Geographic  Magazine.) 

States  now  leads  the  world  in  coal  production,  furnishing  about  one 
third  of  the  total  supply.  England  for  many  years  headed  the  list, 
and  even  now  furnishes  a  large  amount,  for  its  deposits  are  extensive. 

Charcoal  is  a  variety  of  amorphous  carbon  obtained  by 
heating  wood,  bones,  ivory,  and  other  organic  matter  in 
closed  vessels,  or  by  partially  burning  them  in  the  air. 
There  are  several  varieties.  Wood  charcoal  is  a  black, 
brittle  solid,  and  often  has  the  shape  of  the  wood  from 
which  it  is  made.  It  is  insoluble,  though  its  mineral  im- 
purities can  be  removed  by  acids.  It  burns  without  a  flame 
and  leaves  a  white  ash  (mineral  matter).  The  compact 


CARBON  AND   ITS  OXIDES  —  CARBIDES          271 

varieties  conduct  heat  and  electricity,  but  porous  charcoal 
is  a  poor  conductor.  It  resists  the  action  of  moisture  and 
many  chemicals ;  hence  fence  posts,  telegraph  poles,  and 
wooden  piles  are  often  charred  before  being  put  into  the 
ground.  Most  varieties  are  very  porous,  and  when  thrown 
upon  water  a  lump  of  charcoal  floats  for  some  time,  owing 
to  the  presence  of  air  in  its  pores.  Its  porosity  makes  char- 
coal an  excellent  absorber  of  gases,  some  varieties  absorb- 
ing ninety  times  their  bulk  of  ammonia  gas.  It  also  absorbs 
colored  organic  substances  from  solutions ;  this  is  especially 
true  of  animal  charcoal  (see  below) .  Impure  drinking  water 
may  be  partially  purified  by  charcoal,  which  forms  the 
essential  part  of  many  filters  in  houses.  Charcoal  used  for 
such  a  purpose,  however,  must  be  frequently  renewed  or  often 
heated  to  redness  ;  otherwise  it  becomes  contaminated.  The 
taking  up  of  gases  and  solids  (also  liquids)  by  charcoal  is 
called  adsorption  and  is  ascribed  to  the  adhesion  of  the  sub- 
stances upon  the  very  large  condensing  surface  of  the  porous 
charcoal. 

Besides  the  uses  of  charcoal  mentioned  above,  it  is  used 
as  a  fuel,  in  the  manufacture  of  steel  and  of  gunpowder, 
and  as  a  medicine.  It  reduces  oxides  when  heated  with 
them,  thus  :  — - 

2CuO      +      C      =      2Cu      +      CO2 

Copper  Carbon  Copper  Carbon 

Oxide  Dioxide 

Wood  charcoal  is  made  either  in  a  charcoal  pit  or  kiln,  or  in  a  large 
retort.  Where  wood  is  plentiful,  it  is  loosely  piled  into  the  shape 
shown  in  Figure  35,  and  covered  with  turf  to  prevent  too  free  access 
of  air,  though  small  holes  are  left  at  the  bottom  and  a  larger  one  at 
the  top  (as  a  central  flue),  so  that  sufficient  air  can  pass  through  the 
pile.  The  wood  is  lighted,  and  as  it  slowly  burns  care  is  taken  to 
regulate  the  supply  of  air,  so  that  the  wood  will  smolder  but  not  be 
consumed.  The  volatile  matter  escapes  and  the  charcoal  remains, 
the  average  yield  being  about  20  per  cent  of  the  weight  of  the  wood. 


272  INORGANIC   CHEMISTRY 

This  method  is  crude,  uncertain,  and  wasteful.  Much  charcoal  is 
now  made  by  heating  wood  in  closed  retorts,  no  air  whatever  being 
admitted.  By  this  method,  which  is  called  dry  or  destructive  dis- 
tillation, the  yield  of  charcoal  is  30  per  cent,  and  all  the  volatile  matter 


FIG.  35.  —  Wood  arranged  for  burning  into  charcoal. 

is  saved.  In  the  ordinary  combustion  of  wood,  the  hydrogen  forms 
water  and  the  oxygen  forms  carbon  dioxide ;  but  in  dry  distillation, 
where  no  oxygen  is  present,  much  of  the  hydrogen  forms  volatile 
compounds  with  the  carbon  and  oxygen.  Among  these  volatile 
products  are  methyl  alcohol  (CH4O),  acetone  (CaHeO),  and  acetic 
acid  (C2H4O2).  These  are  commercial  substances,  and  contribute  to 
the  profit  of  the  process.  More  or  less  charcoal  is  obtained  by  heating 
any  compound  of  carbon,  e.g.  sugar  or  starch,  the  charring  being  a 
test  for  carbon.  Such  charcoal,  especially  that  obtained  from  sugar, 
is  very  pure. 

Animal  Charcoal  or  Bone  Black  is  made  by  heating  bones 
in  a  closed  vessel,  or  by  heating  a  mixture  of  blood  and 
sodium  carbonate.  It  contains  only  about  10  per  cent  of 
carbon  distributed  throughout  the  porous  mineral  matter 
(calcium  phosphate)  of  the  bone.  Under  the  name  of  ivory 
black,  animal  charcoal  is  used  as  a  pigment,  especially  in 
making  shoe-blacking.  Bone  black  is  extensively  used  to 
decolorize  sugar  sirups,  oils,  and  other  liquids  colored  by 
organic  matter.  (See  Adsorption,  above.) 


CARBON  AND  ITS  OXIDES  —  CARBIDES          273 

Coke  is  made  by  expelling  the  volatile  matter  from  bitumi- 
nous coal,  somewhat  as  charcoal  is  made  from  wood.  It  is 
left  in  the  retorts  when  coal  is  distilled  in  the  manufacture 
of  illuminating  gas.  On  a  large  scale  it  is  made  by  heating 
a  special  grade  of  bituminous  coal  in  huge  ovens,  often  shaped 
like  a  beehive,  from  which  air  is  excluded  soon  after  com- 
bustion begins.  Sometimes  the  coke  is  made  in  closed 
retorts  constructed  to  save  the  by-products,  —  tar,  am- 
monia (in  the  form  of  ammonium  sulphate  or  ammoniacal 
liquor),  organic  compounds  (such  as  benzene,  phenol,  and 
pyridine),  and  combustible  gases.  This  method  not  only 
yields  more  coke,  but  is  also  more  profitable  because  the 
by-products  are  sold  and  the  combustible  gas  is  used  as  a 
source  of  heat,  light,  and  power.  Coke  is  a  grayish,  porous, 
coherent  solid,  harder  and  heavier  than  charcoal.  It  burns 
with  no  smoke  and  a  feeble  flame.  It  contains  about  90 
per  cent  of  carbon,  the  rest  being  the  mineral  matter  origi- 
nally in  the  coal.  Immense  quantities  of  coke  are  used  in 
the  manufacture  of  iron  and  steel.  It  is  superior  to  coal 
for  this  purpose,  because  it  gives  a  greater  heat  when  burned, 
reduces  oxides  easily,  and  contains  little  or  no  sulphur  or 
other  substances  which  would  impair  the  metallic  product. 
The  calorific  value  of  Connellsville  coke  is  about  7900  calories 
per  gram.  Coke  is  the  fuel  used  in  manufacturing  much  of 
the  pig  iron  in  the  United  States,  and  over  twelve  million 
tons  (or  about  three  fourths  of  the  total  amount)  are  made 
annually  in  the  Connellsville  district,  near  Pittsburg,  Penn- 
sylvania. 

Gas  Carbon  is  the  variety  of  amorphous  carbon  which  is 
gradually  deposited  upon  the  inside  of  the  retorts  during 
the  manufacture  of  illuminating  gas.  It  is  a  black,  heavy, 
hard  solid,  and  is  almost  pure  carbon.  It  is  a  good  con- 
ductor of  electricity,  and  is  utilized  to  some  extent  for  the 


274  INORGANIC  CHEMISTRY 

manufacture  of  the  carbon  electrodes  used  in  electric  lights 
and  electric  batteries. 

Lampblack  is  prepared  by  burning  certain  oils  and  resin- 
ous substances  rich  in  carbon  in  a  limited  supply  of  air. 
The  dense  smoke,  which  is  mainly  finely  divided  carbon,  is 
passed  through  a  series  of  condensing  chambers,  where  it 
is  collected  upon  coarse  cloth  or  a  cold  surface.  Its  forma- 
tion is  illustrated  on  a  small  scale  by  a  smoking  lamp,  the 
deposited  soot  being  substantially  the  same  as  lampblack. 
Lampblack  is  one  of  the  purest  forms  of  amorphous  carbon, 
and  is  used  in  making  printer's  ink  and  certain  black  paints. 

Allotropism.  —  Diamond,  graphite,  and  pure  amorphous 
carbon  (e.g.  sugar  charcoal),  though  exhibiting  essentially 
different  properties,  are  identical  in  composition.  All  are 
carbon.  They  can  be  changed  into  one  another  without 
leaving  a  residue,  the  amorphous  form  into  graphite  and 
finally  into  diamond,  and  the  diamond  itself  into  amorphous 
carbon.  Each  burns  in  oxygen,  and  the  sole  product  is 
carbon  dioxide.  Furthermore,  a  given  weight  of  each 
yields  the  same  weight  of  carbon  dioxide,  e.g.  when  12  gm. 
of  each  are  burned,  44  gm.  of  carbon  dioxide  are  always 
produced.  There  is  no  doubt  about  their  identity.  Ele- 
ments which  exist  in  two  or  more  distinct  varieties  are  called 
allotropic,  and  the  property  of  assuming  more  than  one 
variety  is  called  allotropism  or  allotropy  (from  Greek  words 
meaning  "another  form").  One  variety  is  called  an  allo- 
trope  or  an  allotropic  modification  of  the  other.  Allotropism 
is  believed  by  many  to  be  due  to  a  difference  in  the  number 
of  atoms  in  a  molecule  of  the  element,  and  by  others  to  a 
difference  in  the  amount  of  chemical  energy  in  the  allotrope. 

Chemical  Properties  of  Carbon.  —  Carbon  unites  directly 
with  many  elements,  though  the  temperature  must  usually 


CARBON  AND  ITS  OXIDES  —  CARBIDES          275 

be  raised  to  bring  about  combination.  Fluorine  is  the 
only  element  with  which  it  unites  at  ordinary  temperatures. 
The  formation  of  acetylene  (C2H2)  from  carbon  and  hydrogen 
takes  place  at  the  temperature  of  the  electric  arc.  Carbon 
unites  with  many  metals  and  non-metals  at  the  high  tem- 
peratures produced  in  the  electric  furnace,  thereby  forming 
carbides.  (See  end  of  this  chapter.)  It  also  unites  directly 
with  sulphur  to  form  carbon  disulphide  (CS2)  in  a  special 
form  of  furnace.  A  conspicuous  and  important  chemical 
property  of  carbon  is  its  behavior  with  oxygen.  It  unites 
directly  with  oxygen  at  elevated  temperatures  to  form  car- 
bon dioxide  (CO2)  or  carbon  monoxide  (CO),  depending 
upon  conditions  which  are  discussed  under  these  gases. 
Considerable  heat  is  developed  during  the  uniting  with 
oxygen.  The  use  of  various  forms  of  carbon  and  carbona- 
ceous substances  as  fuel  is  based  on  this  fact.  The  thermal 
equation  for  the  transformation  of  carbon  (in  the  form  of 
charcoal)  into  carbon  dioxide  is  — 

C      +      O2     =     C02     +     97,000  cal. 

Carbon        Oxygen  Carbon 

Dioxide 

Carbon  reduces  oxides,  a  simple  illustration  being  the  re- 
duction of  copper  oxide,  thus :  - 

2CuO     +     C     =     CO2       +      2Cu 

Copper  Carbon         Carbon  Copper 

Oxide  Dioxide 

This  property  is  applied  in  the  reduction  of  ores  of  iron  and 
other  metals. 

OXIDES  OF  CARBON 

Carbon  and  oxygen  do  not  unite  at  the  ordinary  tem- 
perature. But  when  carbon  is  heated  in  air,  in  oxygen, 
or  with  some  oxides,  carbon  dioxide  (CO2)  is  usually  formed; 


276  INORGANIC   CHEMISTRY 

if  the  supply  of  oxygen  is  limited,  carbon  monoxide  (CO) 
is  also  formed. 

Occurrence  and  Formation  of  Carbon  Dioxide.  —  The 
occurrence  of  carbon  dioxide  in  the  atmosphere  (3  to  4 
parts  in  10,000)  and  in  many  natural  waters  (especially 
mineral  waters)  has  already  been  mentioned.  It  is  one 
product  of  ordinary  combustion,  respiration  of  animals, 
fermentation,  and  decay;  in  all  these  processes  the  carbon 
comes  from  organic  matter,  while  the  oxygen  comes  for  the 
most  part  from  the  air,  though  some  is  supplied  by  the 
organic  matter. 

Ordinary  combustion  is  the  chemical  combining  of  carbon 
and  oxygen.  Hence,  when  carbon  or  a  substance  containing 
it  is  burned  in  an  excess  of  air  or  oxygen,  carbon  dioxide  is 
formed.  The  equation  for  this  chemical  change  is  — 

C     +     02  CO2 

Carbon        Oxygen  Carbon 

Dioxide 

Carbon  dioxide  is  formed,  therefore,  by  the  combustion  of 
such  common  carbonaceous  substances  as  wood,  coal, 
charcoal,  coke,  oils,  waxes,  cotton,  bone,  starch,  sugar,  meat, 
bread,  alcohol,  camphor,  and  illuminating  gas. 

The  continuous  oxidation  of  the  tissues  of  the  body  pro- 
duces carbon  dioxide.  (See  Relation  of  Oxygen  to  Life.) 
And  if  we  exhale  the  breath  through  a  glass  tube  into  calcium 
hydroxide  (limewater),  the  carbon  dioxide  which  is  in  the 
breath  turns  the  limewater  cloudy  or  turbid  —  the  usual 
test  for  carbon  dioxide.  The  equation  for  the  change  is  - 

CO2     +     Ca(OH)2     =     CaCO3     +     H2O 

Carbon  Calcium  Calcium  Water 

Dioxide  Hydroxide  Carbonate 

When  vegetable  and  animal  matter  decays,  carbon  dioxide 
is  one  product.  Many  kinds  of  organic  matter  ferment, 


CARBON  AND  ITS  OXIDES  —  CARBIDES          277 

especially  those  containing  certain  kinds  of  sugar.  By 
alcoholic  fermentation  the  sugar  changes  into  carbon  dioxide 
and  alcohol  (see  Alcohol),  thus:  — 

C6H12O6     =     2CO2      +      2C2H6O 

Sugar  Carbon  Alcohol 

Dioxide 

The  Preparation  of  Carbon  Dioxide  is  accomplished  in 
the  laboratory  by  the  interaction  of  a  carbonate  and  an  acid. 
Calcium  carbonate  (limestone  or  marble)  and  dilute  hydro- 
chloric acid  are  usually  used.  The  equation  for  the  chemical 
change  is  - 

CaCO3      +       2HC1       =        CO2       +      CaCl2     +     H20 

Calcium  Hydrochloric  Carbon  Calcium  Water 

Carbonate  Acid  Dioxide  Chloride 

This  gas  may  also  be  prepared  by  burning  matter  contain- 
ing carbon  and  by  strongly  heating  carbonates  (as  in  making 
lime),  thus:  - 

C6H12O6   +    6O2       =       6CO2       +       6H2O 

Sugar  Oxygen  Carbon  Water 

Dioxide 

CaCO3  =         C02        +         CaO 

Calcium  Carbon  Lime 

Carbonate  Dioxide 

Properties  of  Carbon  Dioxide.  —  This  gas  has  many  im- 
portant properties  besides  those  mentioned  under  The 
Atmosphere.  It  has  a  slight  taste  and  odor,  but  no  color. 
It  is  one  and  a  half  times  heavier  than  air,  and  a  liter  under 
standard  conditions  weighs  1.977  gm.  On  account  of  its 
weight  it  is  usually  collected  by  the  displacement  of  air 
(like  chlorine  and  hydrochloric  acid  gas).  Its  weight  is 
also  one  reason  why  it  collects  at  the  bottom  of  abandoned 
or  deep  wells,  in  some  valleys  near  volcanoes,  and  in  mines 
after  explosions.  '  At  the  ordinary  temperature  and  pressure, 
water  dissolves  its  own  volume  of  carbon  dioxide.  Under 


278  INORGANIC   CHEMISTRY 

increased  pressure  more  gas  dissolves,  which  escapes  to  some 
extent  when  the  pressure  is  removed.  (See  Carbonic  Acid, 
below.)  Hence  "  soda  water,"  which  is  made  by  forcing 
carbon  dioxide  into  water,  effervesces  and  froths  when  drawn 
from  the  soda  fountain.  Many  natural  waters  and  manu- 
factured beverages  (such  as  champagne)  sparkle  and  ef- 
fervesce for  the  same  reason.  The  critical  temperature 
is  31°  C.  and  the  critical  pressure  is  about  72  atmospheres. 
Hence  this  gas  can  be  readily  liquefied.  It  was  first  liquefied 
by  Faraday  by  the  method  similar  to  that  used  for  chlorine. 
Liquid  carbon  dioxide  is  now  made  in  large  quantities  by 
forcing  the  gas  into  steel  cylinders  (by  powerful  pumps)  which 
are  cooled  by  water  during  the  operation;  the  gas  is  some- 
times obtained  from  the  fermenting  vats  of  breweries,  though 
often  prepared  by  decomposing  magnesium  or  calcium  car- 
bonate. When  liquid  carbon  dioxide  is  allowed  to  escape 
into  the  air,  a  portion  evaporates  quickly  and  thereby  with- 
draws heat  from  the  remainder;  if  sufficiently  cooled,  it 
becomes  white,  snowlike,  solid  carbon  dioxide.  The  latter 
is  used  to  produce  low  temperatures ;  a  paste  of  ether  and 
solid  carbon  dioxide  has  a  temperature  of  —  80°  C.,  and  in 
a  vacuum  it  may  fall  to  —  100°  C. 

Carbon  dioxide  does  not  burn,  but  extinguishes  many 
burning  substances,  such  as  a  blazing  stick  or  lighted  candle, 
the  latter  being  used  occasionally  to  detect  carbon  dioxide 
in  wells,  caves,  and  mines.  Air  containing  from  2.5  to  4 
per  cent  of  carbon  dioxide  will  extinguish  small  flames. 
Hence  the  gas  is  sometimes  used  to  extinguish  fires.  A  stream 
of  gas  forced  upon  a  small  blaze  will  often  prevent  a  serious 
fire.  In  the  portable  fire  extinguishers  and  chemical  engines, 
the  carbon  dioxide  is  generated  rapidly  by  the  interaction 
of  sulphuric  acid  and  dissolved  sodium  bicarbonate  (HNaCO3), 
and  the  pressure  of  the  generated  gas  forces  the  saturated 
solution  of  carbon  dioxide  out  of  the  extinguisher. 


CARBON  AND   ITS  OXIDES  —  CARBIDES          279 

Carbon  dioxide  is  reduced  by  heated  carbon  into  carbon 
monoxide, .  thus :  — 

C02     +      C      =      2  CO 

Carbon  Carbon  Carbon 

Dioxide  Monoxide 

Several  metals  also  decompose  carbon  dioxide.  With 
moderately  heated  zinc  the  chemical  change  is  expressed 
thus : — 

CO2        +        Zn      =      ZnO          +          CO 

Carbon  Zinc  Zinc  Carbon 

Dioxide  Oxide  Monoxide 

With  burning  magnesium,  sodium,  or  potassium  the  chemical 
change  is  as  follows :  — 

CO2      +      2Mg       =       2MgO      -f      C 

Carbon  Magnesium  Magnesium  Carbon 

Dioxide  Oxide 

Carbon  dioxide  combines  directly  with  several  oxides,  thus :  — 
CO2      +      CaO       =      CaCO3 

Carbon  Calcium  Calcium 

Dioxide  Oxide  Carbonate 

It  also  unites  with  water  to  form  carbonic  acid  (H2C03) 
(see  below). 

Relation  of  Carbon  Dioxide  to  Life.  —  Animals  die  when 
put  into  carbon  dioxide.  It  cuts  off  the  supply  of  oxygen 
just  as  water  does  from  a  drowning  man.  The  presence 
of  a  small  quantity  in  the  air  is  objectionable.  On  the  other 
hand,  carbon  dioxide  is  an  essential  food  of  plants.  Through 
their  leaves  and  other  green  parts  they  absorb  carbon 
dioxide,  from  the  atmosphere,  decompose  it,  reject  part  of 
the  oxygen,  and  store  up  the  carbon  as  starch  ((C6HioO5)n). 
The  sunlight  and  the  green  matter  (chlorophyll)  aid  the  plant 
in  making  its  food  out  of  the  water  (obtained  through 
the  roots  from  .the  soil)  and  the  carbon  of  the  carbon  dioxide 
obtained  from  air.  Plants  in  this  way  help  keep  the  atmos- 
phere free  from  an  excess  of  carbon  dioxide. 


280  INORGANIC  CHEMISTRY 

Carbonic  Acid.  —  Carbon  dioxide  gas  is  often  called 
carbonic  acid  gas,  or  simply  carbonic  acid.  The  latter  term 
when  applied  to  carbon  dioxide  (CO2)  is  incorrect.  When 
carbon  dioxide  is  passed  into  water,  it  combines  to  a  slight 
extent  with  the  water  and  forms  a  weak,  unstable  acid, 
which  is,  strictly  speaking,  carbonic  acid.  The  equation 
for  this  change  is  - 

CO2        +        H2O      =      H2CO3 

Carbon  Water  Carbonic 

Dioxide  Acid 

Such  a  solution  reddens  blue  litmus  and  decolorizes  pink 
phenolphthalein,  though  its  action  on  these  indicators  is 
rather  feeble.  In  terms  of  the  theory  of  electrolytic  dis- 
sociation it  is  a  weak  acid,  i.e.  it  ionizes  only  to  a  small 
degree,  the  ions  being  H+  and  HCO3~  (.17  per  cent  in  tenth 
normal  solution  at  18°  C.).  Carbonic  acid  is  unstable  and 
easily  breaks  up  by  gentle  heat  into  carbon  dioxide  and 
water,  thus :  — 

H2CO3  =  CO2  +  H2O 

Carbon  dioxide  is  sometimes  called  carbonic  anhydride,  to 
denote  its  relation  to  the  acid.  Carbonic  acid  is  dibasic 
and  forms  two  classes  of  salts  —  normal  and  acid  carbonates. 

Carbonates  are  salts  corresponding  to  carbonic  acid.  They 
are  very  common  substances.  The  most  abundant  natural 
carbonates  are  those  of  calcium,  magnesium,  and  iron. 
Immense  quantities  of  sodium  and  potassium  carbonates 
are  manufactured. 

A  few  carbonates  are  formed  by  direct  combination  of  an 
oxide  and  carbon  dioxide,  but  many  are  formed  by  the  in- 
teraction of  carbon  dioxide  and  a  base,  thus :  — 

CO2      +      Ca(OH)2      =       CaCO3      +      H20 

Calcium  Calcium 

Hydroxide  Carbonate 


CARBON  AND   ITS  OXIDES  —  CARBIDES          281 

Others  are  formed  by  the  interaction  of  a  soluble  carbonate 
and  another  soluble  salt,  and  are  readily  produced,  since  all 
except  those  of  sodium,  potassium,  and  ammonium,  and 
certain  acid  carbonates  are  practically  insoluble  in  water. 
There  are  two  classes  of  carbonates,  the  normal  and  the 
acid.  Normal  sodium  carbonate  is  Na2CO8,  and  acid  sodium 
carbonate  is  HNaCO3.  The  latter  is  called  sodium  bicar- 
bonate. Normal  calcium  carbonate  is  CaCO3,  and  cal- 
cium bicarbonate  or  acid  calcium  carbonate  is  H2Ca(CO3)2; 
the  latter  salt  is  unstable,  and  is  decomposed  by  gentle 
heat  into  normal  calcium  carbonate.  The  acid  calcium 
carbonate  is  formed  from  the  normal' carbonate  by  an  excess 
of  carbon  dioxide.  When  carbon  dioxide  is  passed  into 
water  containing  the  insoluble  normal  calcium  carbonate  in 
suspension,  the  soluble  acid  calcium  carbonate  is  formed, 
thus : — 

H2COS    +    CaCO3    =    H2Ca(CO3)2 

Carbonic  Calcium  Acid  Calcium 

Acid  Carbonate  Carbonate 

Now  when  this  solution  of  acid  calcium  carbonate  is  heated, 
the  decomposition  takes  place  thus:  — 

H2Ca(CO3)2  =   CaC03      +      CO2     +     H2O 

Acid  Calcium  Calcium  Carbon  Water 

Carbonate  Carbonate  Dioxide 

Since  many  underground  waters  contain  carbon  dioxide, 
these  waters  dissolve  the  limestone  (CaCO8)  over  which  they 
pass,  forming  "  hard  "  water.  When  the  dissolved  acid 
calcium  carbonate  is  decomposed  by  heat  or  in  some  other 
way,  the  calcium  carbonate  is  reprecipitated.  (See  Stalac- 
tites (under  Calcium  Carbonate),  Natural  Waters,  and 
Hardness  of  Water.) 

Composition  of  Carbon  Dioxide.  —  If  a  known  weight  of  pure 
carbon,  such  as  diamond  or  graphite,  is  burned  in  oxygen,  it  is  found 
that  for  12  parts  of  carbon  used  there  are  44  parts  of  carbon  dioxide 
formed.  Hence  12  parts  of  carbon,  unite  with  32  parts  of  oxygen. 


282 


INORGANIC   CHEMISTRY 


Since  the  vapor  density  of  the  gas  is  1.375  (if  Oz  —  32),  the  molecular 
weight  must  be  44.  These  facts  necessitate  the  formula  CO2. 

History  of  Carbon  Dioxide.  — •  This  gas  was  described  in  the  seven- 
teenth century  by  Van  Helmont,  who  called  it  gas  sylvestre.  He  pre- 
pared it  by  the  interaction  of  acids  a'nd  carbonates,  detected  it  in 
mineral  water,  and  observed  its  formation  during  combustion  and 
fermentation,  as  well  as  its  action  on  animals  and  flames.  Black,  in 
1755,  showed  that  carbon  dioxide  is  essentially  different  from  ordinary 
air,  and  that  the  gas  is  readily  obtained  from  magnesium  and  calcium 
carbonates.  Since  the  gas  was  combined  or  "fixed"  in  these  sub- 
stances, he  called  the  gas  fixed  air.  His  work  was  verified  in  1774  by 
Bergman,  who  called  the  gas  acid  of  air.  Lavoisier  first  proved  it 
to  be  an  oxide  of  carbon. 

Carbon  Monoxide  is  formed  when  carbon  is  burned  in  a 
limited  supply  of  air,  thus:  — 

20       +        O2         =          2  CO 

Carbon  Oxygen  Carbon  Monoxide 

If  carbon  dioxide  is  passed  over  heated  charcoal,  the  prod- 
uct is  carbon  monoxide.  That  is,  carbon  reduces  carbon 
dioxide  to  carbon  monoxide,  the  equation  for  the  change 
being  — 

C02  +         C  =  2  CO 

Carbon  Dioxide  Carbon  Carbon  Monoxide 

These  chemical  changes  take  place  in  every  coal  fire  (Fig. 
36).  The  oxygen  (of  the  air)  entering  at  the  bottom  of  the 


2  CO  +  O2  =  2  CO2 
C02  +  C  =  2  CO 
2  C  +  2  02  =  2  CO2 


JL 


FIG.  36.  —  Essential  chemical  changes  during  combustion  In  a  coal  fire. 


CARBON  AND  ITS  OXIDES  —  CARBIDES          283 

fire  unites  with  the  carbon  to  form  carbon  dioxide ;  the  latter 
gas  in  passing  through  the  hot  carbon  of  the  fire  is  reduced 
to  carbon  monoxide.  Some  of  the  carbon  monoxide  escapes, 
while  some  burns  with  a  flickering  bluish  flame  on  the  top 
of  the  fire,  and  forms  carbon  dioxide. 

If  steam  is  passed  through  a  hard  coal  or  coke  fire,  carbon 
monoxide  and  hydrogen  are  formed;  this  mixture  enriched 
by  oil  vapors  is  called  water  gas  and  is  used  as  an  illuminant 
(page  299).  If  steam  and  air  are  passed  through  a  carbon 
fire,  the  gaseous  product  also  contains  nitrogen  and  carbon 
dioxide ;  this  mixture  is  called  producer  gas  and  is  used  as  a 
fuel  in  industrial  processes. 

Preparation  of  Carbon  Monoxide.  —  This  gas  is  usually 
prepared  in  the  laboratory  by  gently  heating  a  mixture  of 
oxalic  acid  and  sulphuric  acid,  and  collecting  the  gaseous 
product  over  water.  The  oxalic  acid  decomposes  thus :  — 

C2H2O4  CO  +          C02       +        H2O 

Oxalic  Carbon  Carbon  Water 

Acid  Monoxide  Dioxide 

The  carbon  dioxide  can  be  removed  by  passing  the  mixed 
gases  through  a  solution  of  sodium  hydroxide. 

Properties  of  Carbon  Monoxide.  —  It  is  a  gas  without 
color,  odor,  or  taste,  and  is  only  slightly  soluble  in  water. 
It  burns  with  a  bluish  flame,  forming  carbon  dioxide,  thus  :  — 

2  CO       +       Oa     =      2CO2 

Carbon  Carbon 

Monoxide  Dioxide 

Carbon  monoxide  is  extremely  poisonous,  and  it  is  very 
dangerous  because  the  lack  of  odor  prevents  its  detection 
in  time  to  escape  its  stupefying  effect.  Many  deaths  have 
been  caused  by  breathing  air  containing  it.  Carbon  mo- 
noxide forms  a  rather  stable  compound  with  the  haemoglobin 
of  the  blood,  and  persons  who  have  been  poisoned  by  it 


284  INORGANIC  CHEMISTRY 

cannot  usually  be  revived  by  air,  as  in  the  case  of  suffoca- 
tion by  carbon  dioxide.  It  is  an  ingredient  of  water  gas 
and  most  ordinary  illuminating  gas.  Care  should  always 
be  taken  to  prevent  the  escape  of  illuminating  gas  (as  well 
as  the  gas  from  a  coal  stove  or  furnace)  into  occupied  rooms. 
At  a  high  temperature  carbon  monoxide  readily  removes 
combined  oxygen,  and  is,  therefore,  an  important  reducing 
agent,  e.g.  in  the  manufacture  of  iron  from  iron  ores  in  the 
blast  furnace.  This  chemical  change  may  be  represented 
thus : — 

Fe203      +       SCO      =      2Fe      +      3CO2 

Iron  Carbon  Iron  Carbon 

Oxide  Monoxide  Dioxide 

Carbon  monoxide,  which  is  sometimes  called  carbonic 
oxide,  forms  no  acid,  and  therefore  no  salts.  It  does  not 
turn  calcium  hydroxide  (limewater)  milky,  thus  being 
readily  distinguished  from  carbon  dioxide.  Its  blue  flame 
distinguishes  it  from  all  other  gases  which  burn.  It  unites 
directly  with  chlorine,  to  form  carbonyl  chloride  (phosgene, 
COC12),  and  with  some  metals,  forming  metallic  carbonyls, 
e.g.  nickel  carbonyl  (Ni(CO)4). 

CARBIDES 

Carbides  are  compounds  of  carbon  and  certain  elements, 
especially  metals.  They  are  numerous,  owing  to  their  ready 
formation  when  carbon  and  oxides  are  heated  in  an  electric 
furnace.  Calcium  carbide  and  silicon  carbide  are  impor- 
tant commercial  substances.  Several  carbides  yield  com- 
pounds of  carbon  and  hydrogen  by  interaction  with  water, 
acetylene  (C2H2)  and  methane  (CH4)  being  commonly 
formed. 

Calcium  carbide,  CaC2,  is  made  on  a  large  scale  by  heating 
a  mixture  of  lime  and  coke  (a  form  of  carbon)  in  an  electric 


CARBON  AND  ITS  OXIDES  —  CARBIDES 


285 


furnace.       The     chemical    change     may     be     represented 
thus : — 

3C       +       CaO       =       CaC2       +       CO 

Carbon  Lime 


=       CaC2 

Calcium 
Carbide 


Carbon 
Monoxide 


The  furnace  in  which  calcium  carbide  is  made  is  sketched 
in  Fig.  36  a.  The  mixture  of  coke  and  lime  (shown  in  the 
furnace)  is  intro- 
duced through  the 
trap  cover  A  and 
slowly  sinks  down 
into  the  space  where 
the  intense  heat  is 
produced  by  the  elec- 
tricity as  it  passes 
between  the  elec- 
trodes G  and  E,  E. 
The  liquid  calcium 
carbide  is  drawn  off 
through  F.  The  car- 
bon monoxide  rises 
through  the  pipes  D, 
D  and  enters  the  up- 
per part  of  the  furnace  together  with  air  supplied  through 
C,  C;  this  mixture  burns  and  heats  the  coke  and  lime. 
The  waste  gases  escape  through  B. 

Calcium  carbide  is  a  hard,  brittle,  dark  gray,  crystalline 
solid.  Its  specific  gravity  is  about  2.2.  With  water  calcium 
carbide  forms  acetylene  (page  290).  Owing  to  its  vigorous 
and  rapid  reaction  with  water,  calcium  carbide  is  packed  and 
sold  in  air-tight  cans. 

Silicon  carbide,  SiC,  is  a  very  hard,  dark-colored  crystalline 
solid  prepared  by  heating  coke  arid  sand  (silicon  dioxide) 
in  an  electric  furnace.  (See  Carborundum.) 


FIG.  36  a. 


Electric  furnace  for  making  cal- 
cium carbide. 


286  INORGANIC  CHEMISTRY 


PROBLEMS  AND  EXERCISES 

1.  The  specific  gravity  of  charcoal  is  about  1.5.     Why  does  it 
float  on  water? 

2.  How  many  grams  of  calcium  carbonate  are  needed  to  pre- 
pare 132  gm.  of  carbon  dioxide? 

3.  Ten  tons  of  coke  were  burned,  and  only  35  tons  of  carbon 
dioxide  were  produced.     Calculate  the  per  cent  of  carbon  in  the 
coke. 

4.  Ten  grams  of  carbon  dioxide  were  prepared  by  the  inter- 
action of  marble  and  hydrochloric  acid.     How  many  cubic  centi- 
meters of  marble  (specific  gravity  2.65)  were  used? 

5.  Suppose  17  gm.  of  carbon  are  completely  and  freely  burned 
in  air  containing  21  per  cent  of  oxygen  by  volume.     Calculate 
(a)  the  volume  of  air  needed,  and  (6)  the  volume  and  the  weight 
of  the  product. 

6.  A  mass  of  limestone  is  completely  transformed  by  heat  into 
a  solid  and  a  gas  —  the  usual  products.     The  gas  measures  2000  1. 
(standard   conditions).     Calculate   the  weight   of   the   three   sub- 
stances involved  in  the  reaction. 

7.  Calculate  the  simplest  formulas  from  the  following  data: 
(a)  C  =  74.07,  N  =  17.29,  H  =  8.64  ;  (6)  C  =  20,  O  =  26.6,  S  =  53.3  ; 

(c)  3  1.  of  an  oxide  of  carbon  weigh  3.78  gm.    (approximately) ; 

(d)  C  =  27.27,  O  =  72.72  ;    (e]  an  oxide  of  carbon  contains  42.857 
per  cent  of  carbon. 

8.  A  cylindrical  tank  holds  250  gm.  of  oxygen.     What  weight 
and  what  volume  of  (a)  carbon  dioxide  and  (6)  carbon  monoxide 
will  it  hold?     (Standard  conditions.) 

9.  Carbon  dioxide  is  heated  with  40  gm.  of  carbon.     What  is 
(i)  the  weight  and   (6)  the  volume  of  the  product?     (Standard 
conditions.) 

10.  Suppose  75  gm.  of  carbon  dioxide  are  passed  slowly  over  hot 
carbon.     Calculate  (a)  the  weight  of  carbon  used  and  (6)  the  vol- 
ume of  the  gaseous  product.      (Standard  conditions.) 

11.  Ten  grams  of  oxalic  acid  are  decomposed.     What  weight  of 
carbon  monoxide  is  formed?     What  volume  of  carbon  monoxide 
at  21°  C.  and  762  mm.? 

12.  What  weight  of  carbon  (97  per  cent  pure)  is  needed  to  reduce 
60  gm.  of  carbon  dioxide  to  carbon  monoxide  ?     What  volume  of  air 
(containing  21  per  cent  of  oxygen  by  volume)  at  19°  C.  and  758  mm. 
is  needed  to  change  the  carbon  monoxide  to  carbon  dioxide? 


CARBON  AND  ITS  OXIDES  —  CARBIDES         287 

13.  If  one  volume  of  carbon  monoxide  and  two  volumes  of  oxy- 
gen are  mixed  and  exploded  in  a  closed  space,  what  will  be  the  vol- 
ume of  the  resulting  gas  or  gases  at  the  original  temperature  and 
pressure  ? 

14.  A  ton  of  calcium  carbide  is  needed.     What  weight  of  lime 
and  coke  must  be  used  ? 

15.  Write  the  formulas  of  the  normal  and  the  acid  carbonates 
of  Ca,  copper,  Pb,  potassium,  Ra,  silver,  Sr,  zinc.     (Use  Valence 
Tables.) 

16.  Starting  with  carbon,  how  would  you  prepare  successively 
(a)  carbon  dioxide,    (6)  calcium  carbonate,    (c)  lime,    (d)  calcium 
hydroxide,  (e)  calcium  chloride? 

17.  Suggest   a  method   of    (a)   obtaining   carbon   from  carbon 
dioxide,  (6)  showing  that  coal  contains  mineral  matter,  and  (c)  test- 
ing a  cave  or  abandoned  mine  for  carbon  dioxide. 

18.  State  and  explain  the  various  chemical  changes  which  occur 
from  the  entrance  of  oxygen  (in  the  air)  below  the  grate  of  a  red- 
hot  coal  fire  to  the  end  of  the  burning  of  the  carbon  monoxide  at 
the  top  of  the  coal. 

19.  Express  the  following  reactions  by  equations  :  (a)  Potassium 
hydroxide  and  carbon  dioxide  form  potassium  carbonate  and  water, 
(6)  barium  hydroxide  and  carbon  dioxide  form  barium  carbonate  and 
water. 

20.  Complete  and  balance  the  following  equations  :    (a)  BaC03 
+  -  =  BaCl2  +  -    -  +  H2O  ;     (6)    HNaCO3  = 


C02. 

21.  Express  the  following  reactions  by  volumetric  equations: 
(a)  Carbon  monoxide  and  oxygen  form  carbon  dioxide  ;    (6)  carbon 
and  water  (vapor)  form  hydrogen  and  carbon  monoxide  ;    (c)  car- 
bon  dioxide  and  carbon  form  carbon  monoxide;     (d)  carbon  and 
oxygen  form  carbon  monoxide  ;   (e)  carbon  and  oxygen  form  carbon 
dioxide. 

22.  Write  the  formulas  of  the  following  compounds  by  apply- 
ing the  principle  of  valence  :    Cadmium  carbonate,  zinc  carbonate, 
barium  carbonate,  ferrous  carbonate.     Calculate  the  weight  and 
the  volume  (standard  conditions)  of  the  gas  liberated  by  the  inter- 
action of  an  acid  and  25  gm.  of  each  carbonate. 


CHAPTER  XVI 

Hydrocarbons  — Methane  —  Ethylene  —  Acetylene  —  Illu- 
minating Gas  —  Flame  —  Bunsen  Burner  —  Oxidizing 
and  Reducing  Flames 

HYDROCARBONS  are  compounds  of  carbon  and  hydrogen. 
They  number  about  two  hundred  and  fifty.  Their  properties 
as  individuals  vary  between  wide  limits,  but  as  a  class  the 
hydrocarbons  are  rather  indifferent  in  chemical  behavior. 
They  are  found  in  petroleum  and  its  products  (kerosene, 
naphtha,  lubricating  oils,  paraffin  wax,  etc.),  in  coal  tar,  in 
coal  gas  and  natural  gas,  and  in  some  essential  oils,  such 
as  turpentine.  On  a  large  scale  they  are  prepared  by  the 
distillation  of  petroleum,  wood,  coal,  and  coal  tar.  Indirectly 
the  hydrocarbons  are  the  source  of  many  other  compounds 
of  carbon,  which  are  extensively  used  in  numerous  industries; 
several  of  these  organic  compounds,  as  they  are  called,  will 
be  described  in  the  next  chapter. 

Methane,  CH4,  is  the  simplest  hydrocarbon.  It  is  found 
in  coal  mines,  being  a  gaseous  product  of  the  processes 
which  changed  vegetable  matter  into  coal.  It  is  called  fire 
damp  by  miners.  It  is  also  formed  in  marshy  places  by 
the  decay  of  vegetable  matter  under  water,  and  is  therefore 
often  called  marsh  gas.  It  forms  about  90  per  cent  of 
natural  gas,  and  approximately  35  per  cent  of  the  illuminating 
gas  obtained  by  heating  coal. 

Methane  is  formed  by  catalysis  when  hydrogen  is  passed 
over  a  mixture  of  carbon  and  nickel  at  250°  C.  It  is  usually 
prepared  by  heating  a  mixture  of  sodium  acetate,  sodium 

288 


METHANE  AND  ETHYLENE  289 

hydroxide,  and  lime.  It  may  also  be  prepared  by  the  inter- 
action of  aluminium  carbide  and  water,  thus  :  — 

A14C8     +     12  H2O     =     3  CH4     +     4  A1(OH), 

Aluminium  Water  -Methane  Aluminium 

Carbide  Hydroxide 

Methane  has  no  color,  taste,  or  odor.  It  burns  with  a 
pale,  bluish  flame.  A  mixture  of  methane  with  oxygen  or 
air  explodes  violently  when  ignited  by  a  spark  or  flame. 
Terrible  disasters  occur  in  coal  mines  from  this  cause.  The 
products  of  the  explosion  are  carbon  dioxide  and  water,  thus  : 

CH4     +     202     =     C02     +     2H20 

Methane          Oxygen  Carbon  Water 

Dioxide 

The  carbon  dioxide,  called  choke  damp  or  black  damp  by 
the  miners,  often  suffocates  those  who  escape  from  the  ex- 
plosion. A  liter  of  methane  weighs  .717  gm.  (standard 
conditions). 

Ethylene,  C2H4,  or  olefiant  gas,  is  formed  by  the  destructive 
distillation  of  wood  and  coal.  It  is  usually  prepared  in  the 
laboratory  by  heating  a  mixture  of  concentrated  sulphuric 
acid  and  ethyl  alcohol  (i.e.  ordinary  alcohol),  and  collecting 
the  gas  over  water.  The  essential  change  is  represented 
thus : — 

C2H6O     =     C2H4     +     H2O 

Alcohol  Ethylene 

Ethylene  is  a  colorless  gas  and  has  a  pleasant  odor.  Its 
critical  temperature  is  9°  C.  and  its  critical  pressure  is  58 
atmospheres,  so  it  can  be  readily  condensed  to  a  liquid. 
The  latter  by  evaporation  in  a  vacuum  produces  a  tempera- 
ture as  low  as  —  140°  C.  A  liter  of  ethylene  weighs  1.25  gm. 
(standard  conditions).  The  gas  burns  with  a  yellow  flame, 
and  is  one  of  the  essential  illuminating  ingredients  of  coal  gas. 
When  ethylene  burns,  the  complete  combustion  is  represented 
thus : — 


290  INORGANIC  CHEMISTRY 

C2H4     +     3O2     =     2CO2     +     2H2O 

Ethylene  Carbon  Water 

Dioxide 

If  mixed  with  oxygen  in  this  proportion  and  ignited,  the  mix- 
ture explodes.     Ethylene  reduces  potassium  permanganate. 

Acetylene,  C2H2,  is  formed  by  the  direct  union  of  hydro- 
gen arid  carbon  when  an  electric  arc  is  produced  between 
two  carbon  rods  in  hydrogen  gas.  This  method  of  forma- 
tion, though  not  convenient,  is  interesting,  because  very  few 
hydrocarbons  have  as  yet  been  formed  directly  from  their 
elements.  A  small  quantity  is  present  in  coal  gas.  It  is  also 
formed  by  the  incomplete  combustion  of  coal  gas,  e.g.  when 
the  flame  of  a  Bunsen  burner  strikes  back  and  burns  at  the 
base.  (See  Bunsen  Burner,  below.)  Acetylene  is  now  pre- 
pared cheaply  on  a  large  scale  by  the  interaction  of  calcium 
carbide  and  water,  thus  :  — 

CaC2     +     2H2O     =     C2H2     +     Ca(OH)2 

Calcium  Water  Acetylene  Calcium 

Carbide  Hydroxide 

Acetylene  is  a  colorless  gas,  and,  if  impure,  has  an  offen- 
sive odor.  It  is  poisonous  if  breathed  in  large  quantities, 
but  much  less  dangerous  than  gases  like  carbon  monoxide. 
Its  density  is  about  .92.  A  liter  weighs  1.162  gm.  (stand, 
cond.).  Water  at  the  ordinary  temperature  dissolves  its  own 
volume  of  the  gas.  Reliable  tests  show  that  acetylene 
does  not  act  upon  any  common  metal  or  alloy,  though  it 
forms  explosive  compounds  with  salts  of  metals,  especially 
of  copper.  As  a  precaution,  copper  and  brass  are  seldom 
used  in  large  vessels  containing  or  generating  acetylene, 
though  they  might  safely  be  used  in  small  lamps. 

The  critical  pressure  is  about  61  atmospheres  and  the  critical  tem- 
perature is  about  35°  C.,  so  acetylene  can  be  readily  liquefied. 
Cylinders  of  liquid  acetylene  have  exploded,  causing  loss  of  life  and 
destruction  of  property,  and  its  use  in  this  form  has  been  prohibited 


ACETYLENE 


291 


in  some  localities.  Under  ordinary  atmospheric '  pressure  acetylene 
gas  does  not  explode,  but  if  subjected  to  a  pressure  of  two  or  more 
atmospheres,  it  will  explode  by  a  shock  or  when  a  spark  or  flame  is 
brought  near  it.  A  mixture  of  acetylene  and  air,  if  ignited,  explodes. 
The  mixture  to  be  explosive,  however,  must  contain  a  large  per  cent 
of  acetylene  (a  condition  hardly  possible  except  from  sheer  careless- 
ness, because  the  disagreeable  odor  reveals  the  presence  of  the  gas). 
Acetylene  gas,  it  is  evident,  must  be  used  with  the  same  precaution 
as  any  ether  illuminating  gas. 

Acetylene  is  found  by  analysis  to  contain  only  carbon  and  hydro- 
gen combined  in  the  ratio  of  12  to  1  by  weight.  Its  vapor  density  is 
.8125  (if  O2=32).  Therefore  its  molecular  weight  is  26  and  its  for- 
mula is  C2H2.  The  graphic  formula  of  acetylene  is  usually  written 
thus:  H-C  =  C-H. 

Acetylene  as  an  Illuminant.  —  Acetylene  burns  in  the  air 
with  a  luminous,  smoky  flame.  But  when  air  is  mixed 
with  the  gas  as  the  latter 
issues  from  a  small  opening, 
the  mixture  burns  with  a  bril- 
liant, white  flame,  which  does 
not  smoke.  It  is  extensively  \ 
used  as  an  illuminant.  The 
flame  is  almost  like  sunlight, 
hence  by  the  acetylene  flame 

most  colors  appear  the  Same  FIG.  37. —Relative  size  of  acetylene 
as  in  daylight.  It  is  adapted  and  illuminating  gas  flames  giving 

for  taking  photographs,  since 
its  action  closely  resembles 
that  of  the  sun.  It  is  a  dif- 
fusive light,  and  the  flame  is 
much  smaller  than  an  ordinary  gas  flame  of  the  same 
lighting  power  (Fig.  37). 

With  a  proper  burner  the  combustion  of  acetylene  is  complete 
and  may  be  represented  thus  :  — 

2C2H2  +  5O2  =  4CO2  +  2HSO 


the  same  amount  of  light.  The 
acetylene  (smaller)  flame  con- 
sumes only  one  tenth  as  much  gas 
an  hour  as  the  illuminating  gas 
flame.  (One  half  actual  size.) 


292  INORGANIC  CHEMISTRY 

Ordinary  gas  burners  cannot  be  used  for  acetylene.  In  acety- 
lene burners  the  gas  issues  from  two  small  holes  drilled  at  an  angle, 
so  that  the  jets  strike  each  other  and  produce  a  flat  flame  (Fig.  38). 


1 


FIG.  38.  —  Acetylene  flame.  FIG.  39.  —  Acetylene  burner. 

Other  holes,  such  as  B,  B,  B,  B,  permit  air  to  be  drawn  in  mechani- 
cally by  the  acetylene  as  it  rushes  through  the  burner.  The  open- 
ings for  the  mixture  are  so  fine  that  the  flame  cannot  strike  back 
and  cause  an  explosion  (Fig.  39). 

Generation  of  Acetylene.  —  The  ease  with  which  acetylene  is  gen- 
erated can  be  shown  by  putting  a  little  water  in  a  test  tube  and  then 
dropping  in  a  small  lump  of  calcium  carbide.  The  gas  bubbles 
through  the  liquid  ;  after  the  action  has  gone  on  long  enough  to  expel 
the  air,  the  acetylene  may  be  lighted  by  holding  a  burning  match  at 
the  mouth  of  the  tube.  On  a  larger  scale,  the  gas  can  be  generated 
by  putting  the  calcium  carbide  into  a  flask  provided  with  a  dropping 
funnel  and  delivery  tube,  and  allowing  water  to  drop  slowly  upon  the 
carbide ;  the  gas  thus  generated  can  be  collected  in  bottles  over 
water.  There  are  two  classes  of  commercial  generators.  In  one, 
water  is  added  to  the  carbide,  but  in  the  other  the  carbide  drops  into 
the  water.  The  intense  heat  liberated  when  calcium  carbide  inter- 
acts with  water  decomposes  acetylene;  hence,  a  generator  to  be 
effective  and  safe  should  be  constructed  so  that  this  heat  will  be  ab- 
sorbed. The  first  class  of  generators  is  dangerous,  except  when  a 
small  quantity  of  gas  is  desired,  as  on  the  lecture  table  or  in  a  lantern. 
In  the  second  class,  a  small  amount  of  calcium  carbide  drops  automat- 
ically into  a  large  volume  of  water  as  fast  as  the  gas  is  needed,  thus 
insuring  a  pure,  cool  gas,  and  eliminating  the  danger  of  an  explosion. 
A  pound  of  calcium  carbide  yields  about  five  cubic  feet  of  acetylene 
gas.  Acetylene  dissolves  in  acetone,  and  cylinders  in  which  con- 
siderable gas  has  been  dissolved  under  pressure  are  used  to  furnish 
gas  for  automobile  lanterns. 

Acetylene  is  an  endothermic  compound,  and  several  of  its  proper- 
ties described  above  are  due  to  its  endothermal  nature.  It  is  formed 


ACETYLENE  293 

from  its  elements  with  absorption  of  heat.  Being  endothermic,  it 
explodes  when  subjected  to  a  shock  (if  the  gas  is  under  a  pressure  of 
two  or  more  atmospheres),  thus:  — 

C2H2  =  2  C  +  H2  +  53,100  cal. 

The  temperature  of  the  oxy-acetylene  flame  is  nearly  3000°  C. ; 
the  thermal  equation  for  the  chemical  change  is  — 

2  C2H2  +  5  O2  =  4  CO2  +  2  H2O  -f  310,000  cal. 

The  flame  when  produced  in  a  suitable  burner  is  used  in  welding 
metals  and  in  dismantling  metal  structures,  e.g.  bridges,  frames  of 
buildings,  abandoned  battleships,  etc.  The  burner  resembles  an 
oxy hydrogen  burner  (Fig.  4). 

Petroleum  is  the  source  of  numerous  hydrocarbons.  It  is 
an  oily  liquid  obtained  from  the  earth  in  many  parts  of  the 
world.  In  the  United  States  the  chief  localities  are  Ohio,  New 
York,  Pennsylvania,  West  Virginia,  Kentucky,  Indiana, 
Colorado,  Texas,  and  California.  The  immense  deposits  in 
Russia  are  in  the  Baku  district  on  the  Caspian  Sea.  Some 
is  also  found  in  Canada,  India,  Japan,  and  Austria. 

Crude  petroleum  is  an  oily  liquid,  with  an  unpleasant 
odor.  Its  color  varies  from  straw  to  greenish  black,  and 
most  kinds  are  greenish  in  reflected  light.  It  usually 
floats  upon  water.  Its  composition  is  complex,  but  all 
varieties  are  essentially  mixtures  of  several  liquid  and  solid 
hydrocarbons.  American  oils  contain  chiefly  members  of 
the  paraffin  series  (i.e.  methane  series).  Some  varieties 
contain  compounds  of  nitrogen  and  of  sulphur. 

In  some  localities  the  oil  issues  from  the  earth,  but  it  is  usually 
necessary  to  drill  through  rocks  and  insert  a  pipe  into  the  porous 
rock  containing  the  oil.  At  first  the  oil  often  "shoots"  out  of  the 
well  in  tremendous  volumes,  owing  to  the  pressure  of  the  confined 
gas,  but  after  a  time  a  pump  is  needed  to  draw  it  to  the  surface. 
The  oil  is  then  forced  by  powerful  pumps  through  large  pipes  to 
central  points  for  storage  or  for  delivery  to  refineries,  which  are  often 
many  miles  from  the  oil  well.  This  network  of  pipes  in  the  eastern 
United  States  is  over  25,000  miles  long. 


294  INORGANIC  CHEMISTRY 

Some  crude  petroleum  is  used  in  making  water  gas  (see 
below)  and  as  fuel  on  locomotives  and  steamships,  but  most 
of  it  is  separated  into  various  commercial  products.  This 
process  is  called  refining.  The  petroleum  is  cleaned  by  set- 
tling and  filtration,  and  then  distilled  in  huge  iron  vessels. 
The  vapors  are  condensed  as  they  pass  through  coiled  pipes 
immersed  in  cold  water.  Certain  products  are  obtained 
from  the  residue  left  in  the  still.  (See  below.) 

The  different  distillates,  which  are  collected  in  separate  tanks, 
are  further  separated  by  redistillation  and  purified  by  special  treat- 
ment. The  commercial  products  thus  obtained  are  petroleum 
ether  (b.  p.  40-70°  C.),  gasolene  (70-90),  naphtha  (90-120),  ben- 
zine (120-150),  kerosene  (150-300) ;  various  grades  of  these  prod- 
ucts are  distinguished  by  boiling  points  or  by  specific  gravity. 
These  liquids  are  mixtures  of  several  different  hydrocarbons.  They 
are  widely  used  as  solvents,  fuels,  illuminants,  and  in  making  gas. 

Gasolene  is  used  as  a  fuel  in  internal  combustion  engines.  The 
vapor  of  gasolene  burns  readily;  if  the  vapor  is  mixed  with  air 
and  the  mixture  is  ignited  properly,  the  combustion  is  so  rapid  that  it 
is  practically  an  explosion  and  the  suddenly  expanded  gas  exerts 
a  pressure  which  is  converted  by  the  machinery  into  steady  and 
continuous  motion.  Kerosene  is  the  well-known  illuminating  oil. 
It  is  carefully  freed  from  inflammable  liquids  and  gases  and  from 
tarry  matter  and  semisolid  hydrocarbons  by  agitating  it  successively 
with  sulphuric  acid,  sodium  hydroxide,  and  water.  Commercial 
kerosene  must  have  a  legal  flashing  point.  Flashing  point  is  **  the 
temperature  at  which  the  oil  gives  off  sufficient  vapor  to  form  a 
momentary  flash  when  a  small  flame  is  brought  near  its  surface." 
The  legal  flashing  point  varies  in  different  localities  from  44°  C. 
to  68°  C. 

From  the  residuum  left  in  the  still  after  the  last  distillation 
lubricating  oils,  vaseline,  and  paraffin  wax  are  obtained  by  further 
treatment.  Mineral  lubricating  oils  have  largely  replaced  animal 
and  vegetable  oils. "  Vaseline  finds  extensive  use  as  an  ointment. 
Paraffin  wax  is  used  to  make  candles,  to  waterproof  paper,  to  extract 
oils  from  plants  and  flowers  ;  also  as  a  coating  for  many  substances, 
thereby  producing  a  smooth  surface  or  facilitating  slow  combustion 
(as  in  certain  matches).  The  final  residue  left  in  the  retorts  is  coke. 


ILLUMINATING  GAS  295 

Hydrocarbons  are  often  extracted  from  it,  some  is  made  into  electric 
light  carbons,  and  some  is  used  as  a  fuel. 

Natural  Gas  is  a  combustible  gas,  which  issues  from  the 
earth  in  many  places,  Methane  forms  about  90  per  cent 
of  the  mixture.  It  is  used  as  a  fuel  for  heating  houses, 
generating  steam,  and  manufacturing  iron,  steel,  glass, 
brick,  and  pottery.  In  Ohio,  Indiana,  and  other  gas- 
producing  regions  of  the  United  States,  wells,  like  petroleum 
wells,  are  drilled  for  the  escape  of  natural  gas,  which  is 
distributed  to  consumers  through  pipes  similar  to  those 
used  for  illuminating  gas. 

Illuminating  Gas.  —  Besides  acetylene  there  are  other 
kinds  of  illuminating  gas.  Coal  gas  and  water  gas  are  the 
most  common. 

Coal  Gas  is  made  by  distilling  bituminous  coal  and  purify- 
ing the  volatile  product.  The  hydrogen  in  the  coal  passes 
off  partly  as  free  hydrogen,  and  partly  in  combination  with 
carbon  as  hydrocarbons  and  with  nitrogen  as  ammonia. 
The  ammonia,  carbon  dioxide,  and  sulphur  compounds 
are  regarded  as  impurities,  and  are  removed  as  completely 
as  possible  before  the  gas  is  delivered  to  the  consumer.  The 
by-products  are  coke,  gas  carbon,  tar,  and  ammoniacal 
liquor.  The  essential  parts  of  a  coal  gas  plant  are  shown 
in  Figure  40. 

The  coal  is  distilled  in  air-tight  O-shaped  retorts  made  of  fire 
clay.  The  volatile  products  escape  through  a  pipe,  and  bubble 
through  water  into  the  hydraulic  main.  Here  some  of  the  tar  is 
deposited  and  the  ammonium  compounds  are  dissolved  by  the  water 
that  flows  constantly  through  the  main.  This  water  is  kept  at  a 
constant  level  and  acts  as  a  "seal"  to  prevent  the  gas  from  passing 
back  into  the  retorts.  The  ammoniacal  liquor  and  tar  flow  into 
separate  wells.  From  the  hydraulic  main  the  gas,  which  is  hot  and 
impure,  passes  into  the  condenser.  This  is  a  long  series  of  vertical 
iron  pipes,  so  constructed  that  the  gas,  must  pass  through  the  entire 


296 


INORGANIC  CHEMISTRY 


ILLUMINATING   GAS  297 

length  of  the  pipes,  while  the  tar  and  ammoniacal  liquor  flow  into 
the  proper  receptacles.  The  main  object  of  the  condenser  is  to 
cool  the  gas  slowly  and  condense  and  remove  the  tar.  An  exhauster, 
in  most  plants,  draws  or  forces  the  gas  from  the  hydraulic  main 
through  the  condenser  into  the  scrubber  and  onward  through  the 
purifiers  into  the  gas  holder.  The  exhauster  also  reduces  the  pressure 
in  the  retorts  and  regulates  the  pressure  in  the  holder.  (See  below.) 
From  the  exhauster  the  gas  passes  into  the  scrubber.  Its  purpose 
is  to  remove  the  remaining  ammonia,  part  of  the  carbon  dioxide 
and  the  hydrogen  sulphide  gas,  and  the  last  traces  of  tar.  Scrubbers 
vary  in  construction.  One  form  is  a  double  tower  filled  with  wooden 
slats  or  with  trays  covered  with  coke  or  pebbles  over  which  ammo- 
niacal liquor  slowly  trickles  in  the  first  part  and  pure  water  in  the 
second.  The  gas  enters  at  the  bottom,  meets  the  descending  liquid, 
and  is  thoroughly  washed.  Another  form  widely  used  consists  of  a 
cylindrical  vessel  in  which  numerous  wooden  slats  revolve  in  com- 
partments and  dip  into  ammoniacal  liquor  or  water  at  the  bottom. 
The  liquid  forms  a  film  on  the  slats  and  absorbs  the  ammonia  and 
other  gases,  while  the  resulting  solution  mixes  with  the  liquor  at 
the  bottom  and  flows  into  the  proper  well.  Sometimes  a  separate 
tar  extractor  is  connected  with  the  scrubber.  This  is  a  tower  filled 
with  perforated  plates,  which  catch  and  remove  the  tar  mechanically 
as  the  gas  passes  through  into  the  scrubber.  From  the  scrubber 
the  gas  passes  into  the  purifiers.  Their  chief  purpose  is  to  remove 
the  remaining  carbon  dioxide  and  sulphur  compounds.  They  are 
shallow,  rectangular  iron  boxes  provided  with  slat  frames  loosely 
covered  with  lime.  In  some  plants  iron  oxide  is  used  as  the  purify- 
ing material. 

The  purified  gas  next  passes  through  a  large  meter,  which  records 
its  volume,  into  a  gas  holder.  The  holder  is  an  enormous,  cylindrical, 
iron  tank  in  which  the  gas  is  stored.  Weights  and  the  pressure 
from  the  exhauster  so  balance  it  that  it  exerts  just  enough  pressure 
to  force  the  gas  through  the  pipes  to  the  consumer. 

A  ton  of  good  coal  yields  about  10,000  cubic  feet  of  gas,  1400 
pounds  of  coke,  120  pounds  of  tar,  20  gallons  of  ammoniacal  liquor, 
and  a  varying  amount  of  gas  carbon.  The  coke,  which  remains  in 
the  retorts  after  distillation,  is  sold  as  fuel.  The  tar,  or  coal  tar  as 
it  is  often  called,  collected  from  the  hydraulic  main  and  condenser, 
is  a  thick,  black,  foul-smelling  liquid.  Some  is  used  for  preserving 
timber,  making  tarred  paper  and  concrete,  and  as  a  protective  paint. 
Most  of  it  is  now  separated  by  distillation  into  its  more  important 


298 


INORGANIC   CHEMISTRY 


ILLUMINATING    GAS  299 

components,  especially  benzene  (CeHe).  The  ammoniacal  liquor 
from  the  hydraulic  main,  condenser,  and  scrubber  is  the  source  of 
ammonia  and  its  compounds.  Gas  carbon  is  the  hard  deposit  which 
collects  on  the  inside  of  the  retort,  and  is  used  in  the  electrical  in- 
dustries. (See  Gas  Carbon.) 

Water  Gas  is  mainly  a  mixture  of  hydrogen  and  carbon 
monoxide.  It  is  made  by  forcing  steam  through  a  mass  of 
hot  anthracite  coal  and  mixing  the  gaseous  product  with 
hot  gases  obtained  from  oil.  The  essential  parts  of  the 
apparatus  are  shown  in  Figure  41. 

Air  is  forced  through  the  coal  fire  in  the  generator,  and  the  hot 
gases  which  are  produced  pass  down  the  carbureter,  up  into  the 
superheater,  and  escape  through  its  top  into  the  open  air.  This 
operation  lasts  about  four  minutes,  and  is  called  the  "  blow."  It 
heats  the  fire  brick  inside  the  carbureter  and  superheater  intensely 
hot,  air  often  being  forced  in  to  raise  the  temperature.  The  air 
valves  and  the  top  of  the  superheater  are  now  closed,  and  the  "  run  " 
begins,  which  lasts  about  six  minutes.  Steam  is  forced  into  the 
generator  at  the  bottom.  In  passing  through  the  mass  of  incan- 
descent carbon  the  steam  and  carbon  interact  thus :  — 

C       +       H2O       =--      CO       +      H2 

Carbon  Steam  Carbon  Hydrogen 

Monoxide 

This  mixture  of  hydrogen  and  carbon  monoxide  burns  with  a 
feeble  though  hot  flame,  and  is  sometimes  used  as  fuel.  Before  this 
mixture  can  be  used  as  an  illuminating  gas  it  must  be  enriched  with 
gases  which  are  illuminants.  Therefore,  the  mixed  gases  pass  to 
the  top  of  the  carbureter,  where  they  meet  a  spray  of  oil.  And  as 
the  gaseous  mixture  passes  down  the  carbureter  and  up  the  super- 
heater, the  hydrocarbons  of  the  oil  are  transformed  by  the  intense 
heat  into  hydrocarbons  which  do  not  liquefy  when  the  gas  is  cooled. 
The  addition  of  hydrocarbons  is  called  carbureting.  From  the 
superheater  the  water  gas  passes  through  the  purifying  apparatus 
into  a  holder. 

Oil  Gas  is  an  illuminating  gas  made  from  petroleum.  When 
petroleum  is  heated  under  proper  conditions,  complex  reac- 
tions occur  which  consist  largely  in  the  decomposition  of  the 


300 


INORGANIC  CHEMISTRY 


heavier  hydrocarbons  ;  this  processes  called  "  cracking."  The 
gaseous  product  by  suitable  treatment  yields  an  illuminating 
gas  containing  a  larger  proportion  (45  per  cent)  of  illuminants 
than  coal  gas.  Pintsch  gas  is  an  oil  gas. 

Characteristics  of  Illuminating  Gases.  —  Coal  gas  and 
water  gas  have  a  disagreeable  odor.  They  are  mixtures. 
The  following  table  shows  the  average :  — 

COMPOSITION  OF  ILLUMINATING  GASES 


CONSTITUKNTS 

COAL  GAS 

WATER  GAS 

Ethylene  (and  other  illuminants)   .... 

6.0 
34.5 

16.6 
19.8 

490 

32  1 

Carbon  monoxide     . 

7  2 

26  1 

Carbon  dioxide         .               ... 

1  i 

30 

Nitrogen     

3.2 

2.4 

Both  kinds  of  illuminating  gas  may  contain  a  little  oxy- 
gen and  traces  of  ammonia  and  hydrogen  sulphide  gas. 
Nitrogen  and  the  last  portions  of  carbon  dioxide  are  impu- 
rities not  easily  removed.  Methane,  hydrogen,  and  carbon 
monoxide  burn  with  a  feeble  (non-yellow)  flame,  and  are 
often  called  diluents;  they  furnish  heat,  but  no  light. 

The  luminosity  of  illuminating  gas  depends  mainly  upon 
the  presence  of  the  hydrocarbons  that  contain  a  relatively 
large  proportion  of  carbon.  The  most  important  illuminants 
in  coal  gas  and  water  gas  are  ethylene  (and  analogous 
hydrocarbons  in  the  same  series),  acetylene,  and  benzene 
(CeH6).  The  utility  of  an  illuminating  gas  depends  upon  its 
luminosity.  This  property  is  measured  by  a  photometer 
and  is  expressed  in  "  candles,"  or  candle  power.  The 
determination  of  candle  power  is  made  by  comparing  the 
light  produced  by  burning  the  gas  in  a  standard  burner 


FLAME  301 

at  the  rate  of  five  cubic  feet  an  hour  with  the  light  pro- 
duced by  a  standard  wax  candle  burning  at  the  rate  of  120 
grains  (7.77  gm.)  an  hour  or  by  a  standard  flame.  Thus, 
a  gas  flame  20  times  brighter  than  the  standard  has  a  can- 
dle power  of  20.  The  candle  power  of  ordinary  coal  gas  is 
about  17,  and  that  of  water  gas  is  about  25.  Ordinary  illumi- 
nating gas  has  a  candle  power  of  about  20.  The  candle 
power  of  oil  gas  is  50  or  more. 

Flame.  —  Chemical  action,  as  we  have  already  seen,  is 
often  accompanied  by  light.  Sometimes  the  light  is  merely 
a  glow,  as  in  the  case  of  burning  charcoal,  or  a  shower  of 
sparks,  seen  when  iron  burns  in  oxygen,  or  a  mere  flash, 
displayed  by  an  explosion  of  hydrogen  and  oxygen.  But 
when  the  interacting  substances  are  gases  or  vapors  from 
volatilized  liquids  or  solids,  the  chemical  change  is  accom- 
panied by  a  more  or  less  quiet  and  continuous  light.  The 
term  flame  is  commonly  applied  to  such  a  light,  though 
a  flame  is  really  a  series  of  chemical  changes  in  which  the 
gases  interact  at  such  a  temperature  that  light  is  produced. 
Popularly,  a  flame  is  a  gas  burning  in  the  air,  i.e.  the  gas 
or  some  of  its  elementary  constituents  are  combining  with 
the  oxygen  of  the  air.  But  since  the  flame  is  due  primarily 
to  the  chemical  combination  of  the  oxygen  of  the  air  and  the 
gas,  it  is  immaterial  where  the  actual  change  takes  place. 
Ordinarily,  the  gas  burns  in  the  surrounding  air,  but  the 
flame  is  produced  just  as  truly,  though  not  so  conveniently, 
in  a  vessel  of  illuminating  gas,  provided,  of  course,  air  is 
supplied.  This  change  of  atmospheres,  so  to  speak,  is  shown 
by  a  simple  experiment  (Fig.  42).  The  lamp  chimney  B 
is  filled  with  illuminating  gas  through  the  bent  tube  D,  and 
its  escape  is  temporarily  prevented  by  closing  the  opening 
in  the  asbestos  cover  A.  The  gas  is  lighted  at  the  lower 
end  of  the  tube  C,  and  when  the  hole  in  A  is  uncovered. 


302 


INORGANIC   CHEMISTRY 


the  flame  rises  in  C  and  continues  at  the  end  within  the 
chimney  as  long  as  air  is  drawn  up  through  C  and  gas  supplied 
through  D.  The  unconsumed  illuminating 
gas  escapes  through  the  hole  in  A,  and  if 
ignited,  burns  as  shown  in  the  figure.  Chem- 
ically both  flames  are  alike.  The  outer  flame 
is  in  an  atmosphere  of  air,  while  the  inner 
flame  is  in  an  atmosphere  of  illuminating  gas; 
but  both  are  due  to  the  combination  of  oxy- 
gen with  the  elementary  constituents  of  the 
illuminating  gas. 

In  an  illuminating  gas  flame  the  gas  itself 
is  burning  in  the  air.     In  a  lamp  flame  the 
burning  gas  comes  from  the  oil  which  is  drawn 
.  up  through  the  wick  and  then  volatilized  by 
bustion  in  iliu-  the  heat.     Similarly,  in  a   candle  flame  the 
minating    gas  burning   gas    comes  from   the   melted   wax. 
The  flame  produced  by  most  burning  hydro- 
carbons is  luminous  and  has  a  yellowish  white  color. 


FIG    4«_ 


Luminous  Flames.  —  The  luminous  hydrocarbon  flame 
has  several  distinct  parts,  and  the  structure  of  the  flame  is 
essentially  the  same,  whether  produced  by  burning  illuminat- 
ing gas,  kerosene  oil,  or  candle  wax.  The  candle  flame  may 
be  taken  as  the  type.  An  examination  of  the  enlarged 
vertical  section  shown  in  Figure  43  reveals  four  somewhat 
conical  portions.  (1)  Around  the  wick  there  is  a  black 
cone  (A),  filled  with  combustible  gases  formed  from  the 
melted  wax.  They  do  not  burn  because  no  oxygen  is  pres- 
ent. With  a  glass  tube  of  fine  bore  it  is  possible  to  draw 
off  these  gases  from  a  large  flame  and  light  them  at  the 
upper  end  of  the  tube.  (2)  Around  the  lower  part  of  the 
dark  cone  is  a  faint  bluish  cup-shaped  part  (B,  B).  It 
is  the  lower  portion  of  the  exterior  cone  where  complete 


LUMINOUS   FLAMES 


303 


combustion  of  the  gases  occurs,  since  plenty  of  oxygen  from 
the  air  reaches  this  portion.  (3)  Above  the  dark  cone  is 
the  luminous  portion  (C).  It  is  the  largest  and  most  im- 
portant part  of  the  flame.  It  is  usually  spoken  of  as  "the 


£"' 


FIG.  43. 


-Typical  candle 
flame. 


FIG.  44.  —  Paper  charred  by  a  candle 
flame. 


flame."  Combustion  is  incomplete  here,  because  little  or 
no  oxygen  can  pass  through  the  exterior  cone.  The  tem- 
perature is  high,  however,  and  the  hydrocarbons  undergo 
complex  changes.  Acetylene  is  probably  formed.  The 
most  characteristic  change  is  the  liberation  of  small  particles 
of  carbon.  This  liberated  carbon,  heated  to  incandescence 
by  the  burning  gases,  makes  the  flame  luminous.  The  carbon 
glows  but  does  not  burn  up,  because  little  or  no  oxygen  is 
present.  A  piece  of  crayon  or  glass  rod  held  in  this  part 
of  the  flame  is  at  once  coated  with  soot,  which  consists  of 
very  fine  particles  of  carbon.  (4)  The  exterior  cone  (D,  D) 
is  almost  invisible.  Here  the  combustion  is  complete, 
because  the  oxygen  of  the  air  changes  all  the  carbon  into 
carbon  dioxide.  That  this  is  the  hottest  region  of  the  flame 
may  be  easily  shown  by  pressing  a  piece  of  stiff  white  paper 
for  an  instant  down  upon  the  flame  almost  to  the  wick. 


304  INORGANIC   CHEMISTRY 

The  paper  will  be  charred  by  the  hot  outer  portion  of  the 
flame,  as  shown  in  Figure  44. 

These  four  portions  may  be  found  in  all  luminous  hydro- 
carbon flames,  whatever  the  shape.  An  ordinary  gas  flame 
is  flattened  by  forcing  the  gas  through  a  narrow  slit  in  the 
burner  tip,  so  that  the  flame  will  give  more  light.  The  blue 
part  is  easily  seen,  however,  when  the  gas  flame  is  turned 
low  or  looked  at  through  a  small  opening;  the  dark  and  yel- 
low parts  are  always  visible — the  latter  being  intentionally 
enlarged.  The  flat  or  circular  flame  of  an  oil  lamp  likewise 
presents  the  same  characteristics. 

The  gaseous  products  of  the  combustion  of  hydrocarbons 
are  water  vapor  and  carbon  dioxide.  A  bottle  in  which 
a  candle  is  burning  has,  at  first,  a  deposit  of  moisture  on 
the  inside;  and  if  the  candle  is  removed  and  lime  water  added, 
the  presence  of  carbon  dioxide  is  shown  by  the  cloudiness 
of  the  limewater.  The  oxygen  needed  by  the  burning 
hydrocarbons  is  obtained  from  the  air.  If  not  enough  oxy- 
gen is  present,  the  flame  smokes,  i.e.  the  carbon  is  thrown  off 
into  the  air  before  the  particles  are  heated  hot  enough  to 
glow.  All  oil  lamps  are  so  constructed  that  air  enters  the 
burner  below  the  flame.  Large  oil  lamps  have  a  central 
opening  through  which  a  large  volume  of  air  passes  up  inside 
the  circular  flame.  Otherwise  the  lamp  would  burn  with 
a  very  smoky  flame. 

The  luminosity  of  hydrocarbon  flames  is  affected  by  other 
factors  besides  the  presence  of  glowing  carbon.  One  of  these 
is  temperature.  Gases  cooled  before  being  burned  give 
very  poor  light.  A  candle  flame  may  be  cooled  enough  to 
extinguish  it.  Thus,  if  a  coil  of  copper  wire  is  lowered  upon 
a  candle  flame,  the  flame  smokes,  loses  its  yellow  color,  and 
finally  goes  out;  but  if  a  coil  of  hot  wire  is  used,  the  flame 
burns  unchanged.  Gases,  as  well  as  solids  and  liquids,  have 
a  kindling  temperature,  i.e.  a  temperature  to  which  they 


THE   BUNSEN   BURNER  AND   ITS  FLAME        305 

must  be  heated  before  they  "catch  fire."  This  temperature 
differs  with  different  substances.  As  we  lower  the  tem- 
perature of  gases  burning  with  a  luminous  flame,  their  lu- 
minosity decreases,  and  below  their  kindling  point  they 
will  not  burn.  The  density  of  the  gases  in  the  flame  and  of 
the  atmosphere  itself  likewise  modifies  luminosity.  The 
flame  of  a  candle  was  found  by  experiment  to  be  smaller 
on  the  top  of  Mont  Blanc  than  at  the  base. 

Not  all  luminous  flames  are  hydrocarbon  flames.  Thus, 
magnesium  burns  with  a  brilliant  flame.  Its  luminosity 
is  due  to  the  incandescence  of  solid  particles  of  magnesium 
oxide.  Similarly,  the  bright  flame  of  burning  phosphorus 
is  accounted  for  by  the  incandescent  particles  of  solid 
phosphorus  pentoxide.  Most  luminous  flames  contain  solid 
particles,  though  in  a  few  cases  luminosity  is  caused  by  the 
combustion  of  gases  under  pressure,  no  solid  particles  what- 
ever being  produced. 

Non-Luminous  Flames.  —  Not  all  flames  are  luminous. 
The  hydrogen  flame  is  almost  invisible  in  air  and  oxygen, 
but  pale  blue  in  chlorine.  The  flames  of  carbon  monoxide 
and  methane  are  also  a  faint  blue.  The  most  common 
non-luminous  flame  is  the  Bunsen  flame. 

The  Bunsen  Burner  and  its  Flame.  —  When  illuminating 
gas  is  mixed  with  air  before  burning,  and  the  mixture  burned 
in  a  suitable  burner,  a  flame  is  produced  which  is  non- 
luminous  and  very  hot.  The  temperature  of  the  hottest 
part  is  about  1500°  C.  This  flame  deposits  no  carbon, 
since  its  products  are  entirely  gaseous.  Such  a  flame  is 
called  the  Bunsen  flame,  for  it  was  first  produced  in  a  burner 
devised  by  the  German  chemist  Bunsen.  This  burner  is 
constantly  used  in  chemical  laboratories  as  a  source  of  heat, 
and  modified  forms  have  numerous  uses.  One  form,  for 
example,  furnishes  the  heat  in  the  gas  range  used  for  cooking. 


306 


INORGANIC  CHEMISTRY 


O 


The  parts  of  a  typical  Bunsen  burner  are  shown  in  Figure  45. 
The  gas  enters  the  base  and  escapes  through  a  very  small 
opening   into  the  long  tube,  which  screws  down 
over  this  opening.     At  the  lower  end  of  the  long 
tube  there  are  two   holes,  through  which  air  is 
drawn  by  the  gas  as  it  rushes  out  of  the  small 
opening.     The  gas  and  air  mix  as  they  rise  in  the 
tube,  and  this  mixture  of  air  and  gas  burns  at  the 
top  of  the  long  tube.     The  size  of  the  air  holes  at 
the  bottom  of  the  long  tube  may  be  changed  by 
a  movable  ring,  thus  varying  the  volume  of  the 
entering  air.     When  the  holes  are  open,  the  typical 
non-luminous,  hot  Bunsen  flame  is  formed.     The 
combustion  of  the  constituents  of  the  hydrocar- 
bons is  practically  complete.     The  non-luminous 
flame  is  free  from  soot,  therefore  apparatus  heated 
by  this  flame  is  not  blackened.    The  Bunsen  flame 
can  be  made  luminous  by  closing  the  air  holes  or 
by  introducing  fine   particles    into 
the  flame,  —  such  as  charcoal  dust, 
finely  divided  metals,   and  sodium 
compounds. 

It  was  formerly  believed  that  the 
non-luminous  character  of  the  Bun- 
sen  flame  is  solely  due  to  the  complete  combustion  of  the 
carbon  by  the  oxygen  of  the  entering  air.  Recent  experi- 
ments have  shown,  however,  that  the  result  is  p?  :-.tly  due 
to  the  diluting  action  of  the  nitrogen. 

The  gas  burns  at  the  top  of  the  tube,  not  inside,  because 
the  proper  mixture  of  gas  and  air  flows  out  more  quickly 
than  the  flame  can  travel  back  through  the  tube  to  the  small 
exit.  If  the  gas  supply  is  slowly  decreased,  the  flame  be- 
comes smaller,  disappears  with  a  slight  explosion,  and  burns 
at  the  exit  inside  the  tube.  A  sudden  draft  of  air,  too  large 


FIG.  45.  —  Parts  of  a  typical 
Bunsen  burner. 


THE   BUNSEN   BURNER  AND    ITS   FLAME       307 

holes  at  the  lower  end  of  the  tube,  or  too  low  gas  pressure 
also  may  cause  the  flame  to  "  strike  back,"  as  this  action  is 
called.  This  change  is  due  to  the  fact  that  the  tube  con- 
tains an  explosive  mixture  of  air  and  illuminating  gas, 
through  which  the  flame  travels  downward  faster  than  the 
mixture  escapes  from  the  tube.  This  modified  flame  has 
a  pale  color  and  disagreeable  odor,  and  deposits  soot. 

The  Bunsen  flame  has  many  characteristic  properties. 
Its  color  is  bluish,  and  the  different  cones  have  different 
tints.  There  are  really  three  cones :  (1)  the  blue  or  greenish 
inner  one  of  unburned  gases;  (2)  the  very  faint  blue  middle 
one;  (3)  and  the  outer  one,  which  is  a  pale  blue,  and  rep- 
resents the  blue  cone  in  the  candle  flame.  The  middle 
and  outer  cones  are  not  always  easily  distinguished;  so  for 
all  practical  purposes  it  is  convenient  to  divide  the  flame 
into  two  parts,  —  an  inner  cone  of  unburned  gases  and  an 
outer  cone  in  which  all  the  carbon  is  consumed.  Combustible 
gases  may  be  drawn  off  by  a  tube  from  the  inner  cone  and 
ignited.  A  match  laid  for  an  instant  across  the  top  of  the 
tube  is  charred  only  at  the  two  points  where  it  touches 
the  outer  cone;  and  a  sulphur  match  suspended  by  a  pin 
across  the  top  of  an  unlighted  burner  is  not  kindled  until 
several  minutes  after  the 
gas  is  first  lighted. 

A  piece  of  wire  gauze 
pressed  down  upon  the  Bun- 
sen  flame  shows  a  dark  cen- 
tral portion  surrounded  by 
a  luminous  ring.  The  flame  H 

is  beneath    the    gauze,   al- 
though the  gas  passes  freely        FlG<  46'  ~~  Wire  gauze  and  flame- 
through  it  and  escapes.    If  the  gas  is  extinguished  and  then 
relighted  above  the  gauze,  it  will  burn  above  but  not  beneath 
(Fig.  46).     The  gauze  cools  the  gas  below  its  kindling  tem- 


ot  kindled  until 

Mavmjv 


308 


INORGANIC   CHEMISTRY 


perature.  The  miner's  safety  lamp  invented  by  Davy  depends 
upon  this  last  principle.  It  is  an  oil  lamp  surrounded  by  a 
cylinder  of  fine  wire  gauze  (Fig.  47).  When 
taken  into  a  mine  where  there  are  explosive 
gases  (such  as  methane  —  see  Methane) ,  the 
flame  continues  to  burn  inside,  though  its 
size  and  color  change.  The  gas  often  enters 
the  lamp  and  burns  inside,  but  the  flame 
within  does  not  ignite  the  gases  without 
because  the  wire  gauze  keeps  them  cooled 
below  their  kindling  temperature.  Hence  an 
explosion  is  often  prevented.  When  miners 
notice  changes  in  the  lamp  flame,  they  usually 

FIG.     47.  — One    seek   a   safe   place, 
form  of  Davy's 
safety  lamp. 

Oxidizing  and  Reducing  Flames,  —  The  outer 

portion  of  the  Bunsen  flame  is  called  the  oxidizing  flame, 
because  here  the  oxygen  is  freely  given  to  substances.  The 
inner  portion  is  called  the  reducing  flame, 
because  here  the  hydrocarbons  withdraw 
oxygen.  A  sketch  of  the  general  relation 
of  these  flames  is  shown  in  Figure  48. 
A  is  the  most  effective  part  of  the  oxidiz- 
ing flame,  and  B  of  the  reducing  flame. 
At  A  metals  are  oxidized,  and  at  B  oxy- 
gen compounds  are  reduced.  Sometimes 
a  long  tube  with  a  small  opening  at  one 
end,  called  a  blowpipe,  is  used  to  produce 
these  flames.  Another  tube  with  a  flat- 
tened top  is  put  inside  the  burner  to  pro- 
duce a  luminous  flame.  The  tip  of  the 
blowpipe  rests  in  or  near  this  flame,  and 

....  ±,  ,.  ill  FIG.    48.  —  The  oxr- 

if  air  is  gently  and    continuously    blown      dizing  (A)  and  re. 
through  the  blowpipe,  a  long,  slender  flame      ducing  (B)  flames* 


•—A 


----- B 


OXIDIZING  AND  REDUCING  FLAMES  309 

is  produced,  called  a  blowpipe  flame  (Fig.  49).  It  is  like 
the  Bunsen  flame  as  far  as  its  oxidizing  and  reducing  prop- 
erties are  concerned.  The  blowpipe  is 
used  in  the  laboratory  and  by  jewelers 
and  mineralogists.  On  a  large  scale  the 
blowpipe  flame  is  used  to  reduce  or  oxidize 
ores.  (See  Compound  Blowpipe.)  Fm  49  _  Blow"pipe 

The  Bunsen  flame  is  extensively  utilized  flame,  showing  oxi- 
in  producing  the  Welsbach  light.  The  £S«  (g^*1 
non-luminous  flame  heats  an  inverted 
bag  or  mantle  of  oxides  of  thorium  and  cerium,  and  the 
mantle  glows  with  an  intense  light.  The  candle  power 
varies  from  40  to  100.  The  proportion  of  thorium  oxide  to 
cerium  oxide  in  the  mantle  is  99  to  1. 

PROBLEMS 

1.  Calculate  the  weight  of  carbon  in  (a)  32  gm.  of  methane, 
(6)  75  gm.  of  ethylene,  and  (c)   145  gm.  of  acetylene. 

2.  What  volume  and  what  weight  of  oxygen  are  needed  for  the 
complete  combustion  of  (a)  15  gm.  of  methane,  (6)  20  gm.  of  ethy- 
lene, and  (c)  25  gm.  of  acetylene?     (Standard  conditions.) 

3.  How  many  grams  of  potassium  chlorate  are  required  to  fur- 
nish the  oxygen  necessary  to  burn  10  liters  of  methane,  and  how 
many  liters  of  each  of  the  products  will  be  formed?     (Standard 
conditions.) 

4.  A  gas  holder  has  a  maximum  capacity  of  12,000  cubic  meters. 
How  much  calcium  carbide  (92  per  cent  pure)  must  be  used  to  fill 
the  holder  with  acetylene  gas  measured  at  20°  C.  and  757  mm.  ? 

5.  What  weight  of  acetylene  can  be  prepared  from  (a)  a  metric 
ton  of  pure  calcium  carbide  and  (6)  a  pound  of  calcium  carbide 
which  is  90  per  cent  pure? 


CHAPTER  XVII 
Other  Carbon  Compounds 

Introduction.  —  It  was  formerly  believed  that  starch,  sugar, 
and  other  compounds  obtained  from  plants  and  animals 
were  produced  by  the  influence  of  some  mysterious  vital  force. 
Such  compounds  were  called  organic,  because  of  their  con- 
nection with  living  things,  i.e.  with  bodies  having  organs; 
and  they  were  sharply  distinguished  from  inorganic  or 
mineral  compounds  obtained  from  the  earth's  crust.  This 
distinction  prevailed  until  Wohler,  in  1828,  prepared  urea  — 
a  distinct  organic  compound  —  from  inorganic  substances. 
Since  then  the  barrier  between  the  two  classes  of  compounds 
has  been  completely  removed.  It  is  now  believed  that  com- 
pounds of  carbon,  whatever  their  source,  are  subject  to  the 
laws  that  govern  all  other  compounds.  The  terms  organic 
and  inorganic  are  still  used,  though  they  have  lost  their 
original  narrow  meaning.  Carbon  forms  a  vast  number  of 
compounds  which  are  related  to  each  other,  and  which  differ 
from  most  compounds  of  other  elements.  It  is  convenient, 
therefore,  to  distinguish  these  compounds  by  the  term 
organic  and  to  study  them  under  the  comprehensive  title  of 
Organic  Chemistry  or  the  Chemistry  of  Carbon  Compounds. 
Several  organic  compounds  have  already  been  discussed  in 
the  chapter  immediately  preceding.  A  few  typical  com- 
pounds only  can  be  considered  in  the  present  chapter. 

Composition  of  Organic  Compounds.  —  The  number  of 
organic  compounds  is  very  large,  but  they  contain  only  a  few 
elements  —  seldom  more  than  four  or  five.  Hydrocarbons, 

310 


OTHER  CARBON  COMPOUNDS        311 

as  already  stated,  contain  only  carbon  and  hydrogen.  Vege- 
table substances,  typified  by  starch,  sugar,  and  fruit  acids, 
contain  carbon,  hydrogen,  and  oxygen.  Animal  substances, 
like  hair,  albumin,  gelatin,  and  muscle  generally  contain 
nitrogen  as  well  as  carbon,  hydrogen,  and  oxygen;  some  also 
contain  sulphur  or  phosphorus.  Artificial  organic  com- 
pounds, like  dyestuffs  and  medicines,  may  contain  any  ele- 
ment, especially  chlorine,  iodine,  and  certain  metals. 

The  number  and  complexity  of  organic  compounds  are  due 
to  several  facts.  (1)  Atoms  of  carbon,  unlike  those  of  most 
elements,  have  power  to  unite  with  themselves.  (2)  Atoms  of 
different  elements  can  be  introduced  into  carbon  compounds. 
Sometimes  these  atoms  are  simply  added,  sometimes  they 
replace  other  atoms,  thus  producing  an  endless  number  of 
addition  and  substitution  products.  (3)  The  same  number 
of  atoms  may  arrange  themselves  differently,  thereby  pro- 
ducing isomeric  compounds  having  different  properties. 
(4)  Organic  compounds  contain  radicals.  These  radicals 
are  groups  of  atoms  analogous  to  hydroxyl  (OH)  and  ammo- 
nium (NH4),  and  like  these  radicals  they  exist  only  in  com- 
bination. They  act  chemically  like  single  atoms  and  pass 
from  one  compound  to  another  without  decomposition.  The 
radical  C2H5  is  called  ethyl.  It  is  present  in  many  organic 
compounds,  and  its  presence  in  ordinary  alcohol  gives  rise 
to  the  scientific  name,  ethyl  alcohol.  Methyl  (CH3)  is 
another  important  radical,  and  phenyl  (C6H5)  is  especially 
common  in  the  benzene  series  of  organic  compounds. 

Structure  of  Organic  Compounds.  —  An  extensive  study 
of  the  properties  of  organic  compounds  has  revealed  many 
facts  about  their  constitution,  i.e.  the  structure  of  their  mole- 
cules. Relatively  little  is  known  about  the  shape,  size,  etc., 
of  molecules,  but  much  is  known  about  the  grouping  of  atoms 
and  of  radicals  in  the  molecules.  These  facts,  which  are 


312  INORGANIC  CHEMISTRY 

ascertained  by  experiment  and  are  often  too  complex  to  be 
expressed  briefly,  may  be  represented  by  suitable  formulas. 
Thus,  the  ordinary  or  empirical  formula  of  alcohol  is  C2H6O. 
But  this  formula  tells  nothing  about  the  relation  these  atoms 
bear  to  each  other.  Experiment  shows  that  (1)  one  hydro- 
gen atom  acts  differently  from  the  other  five,  and  (2)  one 
hydrogen  atom  is  always  associated  with  the  oxygen  atom  in 
chemical  changes.  Hence,  the  formula  C2H6 .  OH  expresses 
more  fully  these  facts.  Such  a  formula  is  called  a  rational 
or  constitutional  formula.  Sometimes  constitution  is  ex- 
pressed by  a  graphic  or  structural  formula.  Thus,  methane 
and  ethane  have  the  graphic  formulas:  — 

H  H   H 

I  I      I 

H-€— H  H— C— C— H 

I  I      I 

H  H   H 

Methane  Ethane 

In  these  diagrams  the  single  lines  represent  a  valence  of  one 
—  nothing  else,  and  the  number  of  lines  connected  with  each 
atom  is  equal  to  the  valence  of  the  element  in  the  compound. 
The  lines  are  sometimes  called  bonds  or  links,  but  they  are 
not  intended  to  represent  attraction  or  any  other  force.  The 
graphic  formula  of  ethyl  alcohol  is  :  — 

H   H' 
I      I 
H— C— C— O— H 

I      I 
H  H 

This  is  not  an  arbitrary  arrangement;  the  facts  mentioned 
above  necessitate  this  general  configuration. 

Classification    of    Organic    Compounds.  —  Organic    com- 
pounds are  divided  and  subdivided  into  many  classes,  mainly 


OTHER  CARBON  COMPOUNDS  313 

for  the  purposes  of  study.  The  most  common  classes  are : 
(1)  Hydrocarbons;  (2)  Alcohols;  (3)  Aldehydes;  (4)  Ethers; 
(5)  Acids;  (6)  Esters;  (7)  Fats,  glycerin,  and  soap; 
(8)  Carbohydrates;  (9)  Benzene  and  its  derivatives; 
(10)  Cyanogen  and  its  derivates;  (11)  Proteins.  Some 
compounds  are  closely  related  and  belong  to  several  of 
these  groups,  while  a  few  common  ones  are  excluded. 

HYDROCARBONS 

Three  hydrocarbons  (methane,  ethylene,  and  acetylene) 
have  been  considered  in  Chapter  XVI.  (See  also  Benzene, 
below.) 

Three  substitution  products  of  methane  are  chloroform  (CHC13), 
iodoform  (CHI3),  and  carbon  tetrachloride  (CC14).  Chloroform 
is  a  heavy  liquid.  It  is  made  by  treating  alcohol  or  acetone  with 
bleaching  powder,  and  is  used  as  an  anaesthetic.  Iodoform  is 
a  yellow  solid.  It  is  made  by  treating  alcohol  or  acetone  with 
iodine  and  sodium  carbonate,  and  is  used  as  an  antiseptic  dressing 
for  wounds.  Carbon  tetrachloride  is  a  heavy  liquid.  It  is  made 
by  passing  dry  chlorine  into  carbon  disulphide  (in  which  a  little 
iodine  acts  as  a  catalyzer).  It  is  used  to  extract  fatty  substance's 
from  seeds,  bones,  and  wool;  certain  non-inflammable  mixtures 
used  for  cleansing  fabrics  (e.g.  "  carbona  ")  contain  carbon  tetra- 
chloride. When  heated,  it  forms  a  heavy,  non-inflammable  vapor, 
and  hence  it  is  used  in  some  fire  extinguishers  (e.g.  "  pyrene  "). 

ALCOHOLS 

Ordinary  or  ethyl  alcohol  is  the  best-known  member  of 
this  group.  It  is  often  called  simply  alcohol.  There  are 
many  alcohols  analogous  to  ethyl  alcohol,  but  the  only  other 
important  one  is  methyl  alcohol. 

The  alcohols  may  be  regarded  as  hydroxides  of  certain  radicals, 
viz.  ethyl,  methyl,  propyl,  etc.1  For  example,  ethyl  alcohol  is  ethyl 
hydroxide,  and  may  be  considered  as  formed  by  replacing  one  hydro- 

1  The  names  of  these  and  similar  radicals  are  derived  from  the  name  of 
the  corresponding  hydrocarbon.  Thus,  methyl  from  methane,  ethyl  from 
ethane,  propyl  from  propane. 


314  INORGANIC  CHEMISTRY 

gen  atom  of  ethane  (C2H6)  by  one  hydroxyl  group  (OH).  Again, 
alcohols  are  analogous  to  metallic  hydroxides,  in  which  the  metal 
is  replaced  by  a  radical,  thus :  — 

C2H5 .  OH  NaOH 

Ethyl  Hydroxide  Sodium  Hydroxide 

Alcohols  and  metallic  hydroxides  have  some  properties  in  common. 
Thus,  both  form  salts  with  acids.  With  acetic  acid,  sodium  hydrox- 
ide forms  sodium  acetate,  while  alcohol  forms  ethyl  acetate.  (See 
Esters.) 

Methyl  Alcohol,  CH3 .  OH,  is  a  colorless  or  slightly  yellow- 
ish liquid,  much  like  ordinary  alcohol.  It  boils  at  about 
66°  C.,  and  burns  with  a  pale  flame  which  deposits  no  soot. 
Methyl  alcohol  causes  blindness  and  even  death.  It  mixes 
with  water  in  all  proportions.  It  is  cheaper  than  ethyl 
alcohol,  and  is  used  as  a  solvent  for  fats,  oils,  and  shellac, 
and  in  the  manufacture  of  varnishes  and  dyestuffs.  Methyl 
alcohol  is  often  called  wood  alcohol  or  wood  spirit,  because 
it  is  one  of  the  products  obtained  by  the  dry  distillation  of 
wood.  (See  Charcoal.) 

Ethyl  Alcohol,  C2H5 .  OH,  is  a  colorless,  volatile  liquid, 
having  a  burning  taste  and  a  pleasant  odor.  Its  specific 
gravity  is  about  0.8.  It  boils  at  about  78°  C.,  and  freezes 
at  about  —  112°  C.  Alcohol  mixes  with  water  in  all  propor- 
tions. The  commercial  variety  contains  about  95  per  cent 
of  alcohol  by  volume.  Absolute  alcohol  contains  over  99 
per  cent  of  alcohol  and  is  prepared  by  distilling  ordinary 
alcohol  with  lime.  Denatured  alcohol  is  a  mixture  of  100 
parts  ethyl  alcohol,  10  parts  methyl  alcohol,  and  a  small  pro- 
portion of  benzine  or  pyridine  (or  a  similar  mixture).  It 
is  unfit  for  drinking,  largely  on  account  of  the  disagreeable 
taste,  but  is  suitable  for  industrial  uses.  Alcohol  burns  with 
a  hot,  nearly  colorless,  non-smoking  flame,  and  is  often  used  as 
a  source  of  heat.  It  is  an  excellent  solvent  for  gums,  oils, 
and  resins,  and  is  therefore  extensively  used  in  the  manufac- 


OTHER  CARBON  COMPOUNDS  315 

ture  of  varnishes,  essences,  extracts,  tinctures,  perfumes, 
and  medicines.  Many  organic  compounds,  as  ether  and 
chloroform,  are  prepared  from  alcohol.  Some  vinegar  is 
made  from  alcohol.  In  museums  alcohol  is  used  to  preserve 
certain  specimens.  Alcohol  is  manufactured  by  the  fermen- 
tation of  sugars  and  starches. 

Fermentation  is  a  general  term  for  the  chemical  changes 
caused  by  compounds  secreted  by  ferments.  The  latter  are 
minute  living  organisms;  the  compounds  they  secrete  are 
called  enzymes.  The  process  and  essential  products  vary 
with  the  nature  of  the  ferment.  The  important  kinds  of 
fermentation  are  alcoholic,  acetic,  and  lactic ;  and  they  pro- 
duce alcohol,  acetic  acid,  and  lactic  acid.  Alcoholic  fermen- 
tation is  caused  by  the  enzyme  zymase  that  is  secreted  by 
ordinary  yeast.  When  yeast  is  added  to  a  solution  of  glu- 
cose, maltose,  or  any  other  fermentable  sugar,  the  yeast 
plants  multiply  rapidly.  The  changes  are  numerous  and 
complex,  but  the  main  products  resulting  from  the  action  of 
the  enzyme  from  the  yeast  upon  the  sugar  are  alcohol  and 
carbon  dioxide,  thus  :  — 

C6H1206       =       2C2H60       +       2C02 

Dextrose  Alcohol  Carbon  Dioxide 

Commerical  alcohol  is  made  from  starch.  The  starch  is 
changed  into  maltose,  etc.,  by  an  enzyme  called  diastase,  in 
the  malt,  and  the  maltose  is  changed  into  alcohol  and  carbon 
dioxide  by  the  zymase  in  the  yeast.  Wine,  beer,  and  distilled 
liquors  are  essentially  mixtures  of  alcohol  and  water.  They 
differ  mainly  in  their  proportion  of  alcohol.  The  particular 
flavor  is  due  to  small  quantities  of  different  substances  which 
are  intentionally  added,  obtained  from  the  raw  materials,  or 
formed  by  special  processes  of  manufacture.  Beer  contains 
from  3  to  7  per  cent  of  alcohol,  wines  from  6  to  20,  rum, 
brandy,  and  whisky  from  40  to  60  or  more  per  cent. 


316  INORGANIC  CHEMISTRY 

ALDEHYDES  AND  KETONES 

Formaldehyde,  CH2O,  is  a  gas,  but  is  usually  used  in  solu- 
tion. It  has  a  penetrating  odor.  The  commercial  solution 
sold  as  formalin  contains  40  per  cent  of  formaldehyde. 
Formaldehyde  is  used  in  the  manufacture  of  dyestuffs  and 
fuming  nitric  acid,  and  as  a  disinfectant.  When  used  for 
the  last  purpose,  formalin  is  vaporized  in  a  special  apparatus 
and  the  vapor  is  conducted  into  the  infected  room.  It  hard- 
ens tissue  and  is  used  as  a  preservative  in  museums  and 
biological  laboratories.  Other  aldehydes  are  benzaldehyde 
(oil  of  bitter  almonds,  C7H60)  and  vanillin  (CgHgOs) ;  both 
are  used  as  flavors. 

Acetone,  C3H6O,  is  a  colorless  liquid  which  has  an  ethereal 
odor.  It  boils  at  about  56°  C.  and  mixes  in  all  proportions 
with  water,  alcohol,  and  ether.  It  is  used  as  a  solvent  for 
fats,  oils,  and  waxes,  and  in  the  preparation  of  smokeless 
powders  and  certain  organic  compounds.  Acetone  is  one  of 
the  products  obtained  by  the  dry  distillation  of  wood.  (See 
Charcoal.) 

ETHERS 

Ordinary  or  ethyl  ether  is  the  best-known  member  of  this 
group. 

Ethyl  Ether,  C4Hi0O,  is  a  colorless,  volatile  liquid,  with  a 
peculiar,  pleasing  taste  and  odor.  It  boils  at  35°  C.,  and  the 
vapor  is  very  inflammable.  The  liquid  should  never  be 
brought  near  a  flame.  It  is  somewhat  soluble  in  water,  and 
it  also  dissolves  water  to  a  slight  extent.  It  mixes  with 
alcohol  in  all  proportions.  It  is  a  good  solvent  for  waxes, 
fats,  oils,  and  other  organic  compounds.  Its  chief  use  is  as 
an  anaesthetic.  Ether  is  manufactured  by  distilling  a  mix- 
ture of  ethyl  alcohol  and  sulphuric  acid  in  the  proper  propor- 
tions. Hence  the  name,  ethyl  or  sulphuric  ether. 


OTHER   CARBON   COMPOUNDS  317 

ACIDS 

This  large  class  of  compounds  is  divided  into  several  series, 
one  of  the  most  important  of  which  is  the  acetic  or  fatty 
series.  Its  best  known  member  is  acetic  acid;  several  of  the 
higher  members  occur  in  fats  and  oils.  These  acids  are 
closely  related  to  hydrocarbons,  alcohols,  and  aldehydes,  as 
may  be  seen  by  the  following  formulas  :  — 

H  H  H 

I  I  I 


H—  C—  H 
1 
H—  C—  H 

i 

H 

H—  C—  (OH) 
1 
H—  C—  H 
1 
H 

C=O 
1 
H—  C—  H 

1 
H 

0=C-(C 

1 

H—  C—  H 

1 
H 

Ethane  Ethyl  Alcohol  Acetic  Aldehyde         Acetic  Acid 

The    characteristic   radical   of   organic  acids  is  COOH  (or 
O=C — O — H),  and  is  called  carboxyl. 

Acetic  Acid,  C2H4O2  or  CH3 .  COOH.  This  is  the  most 
common  organic  acid.  It  is  manufactured  on  a  large  scale 
by  the  dry  distillation  of  wood.  The  dark  red  watery  dis- 
tillate, which  is  called  pyroligneous  acid,  contains  about  10 
per  cent  of  acetic  acid,  besides  methyl  alcohol  and  acetone. 
This  distillate  is  neutralized  with  lime  or  sodium  carbonate, 
and  the  acetate  formed  is  then  decomposed  and  distilled  with 
sulphuric  acid.  The  acetic  acid  which  condenses  in  the  re- 
ceiver may  be  further  purified  by  distilling  it  with  potassium 
dichromate  and  then  filtering  through  charcoal.  Sometimes 
the  pyroligneous  acid  is  distilled  without  neutralizing;  the 
distillate  is  then  dilute,  impure  acetic  acid,  and  is  known  as 
wood  vinegar.  If  sodium  acetate,  prepared  as  described 
above,  is  fused  and  then  distilled  with  concentrated  sul- 
phuric acid,  the  product  is  very  concentrated  acetic  acid. 


318  INORGANIC  CHEMISTRY 

It  is  called  glacial  acetic  acid,  because  at  about  17°  C.  it 
becomes  an  icelike  solid. 

Commercial  acetic  acid  is  a  water  solution  containing 
about  30  per  cent  of  pure  acetic  acid.  It  is  a  colorless  liquid, 
having  a  pleasant  odor  and  a  sharp  taste.  It  is  a  weak  acid, 
a  normal  solution  at  18°  C.  being  dissociated  to  the  extent 
of  about  .4  per  cent  into  the  ions  H+  and  C2H3O2~.  It 
mixes  with  water  and  alcohol  in  all  proportions,  and  like 
alcohol  is  an  excellent  solvent  for  many  organic  substances. 
Recently,  it  has  begun  to  replace  alcohol  as  a  solvent  for 
many  drugs. 

Acetic  acid  is  used  to  prepare  acetates,  dyestuffs,  medi- 
cines, white  lead,  and  in  the  manufacture  of  vinegar. 

Vinegar  is  dilute,  impure  acetic  acid.     It  is  prepared  by  oxidizing 
dilute  alcohol,  the  essential  change  being  represented  thus :  — 
C2H6O     +     O2      =      C2H4O2     +     H2O 

Alcohol  Oxygen  Acetic  Water 

Acid 

The  transformation  is  accomplished  by  fermentation.  (1)  When 
beer,  weak  wines,  or  cider  are  exposed  to  the  air,  they  slowly  become 
sour,  owing  to  the  conversion  of  alcohol  into  acetic  acid.  The  change 
is  caused  by  the  presence  and  activity  of  a  ferment,  known  as  myco- 
derma  aceti,  or  "mother  of  vinegar."  Strong  wines  and  pure  dilute 
alcohol  do  not  become  sour,  because  the  ferment  cannot  live  in 
such  liquids.  (2)  Fruit  juices  and  molasses  contain  fermentable 
sugar  and  ferment  when  exposed  to  the  air  (which  always  contains 
the  necessary  organisms),  forming  alcohol  first  and  finally  vinegar. 
Cider  vinegar  is  made  this  way.  (3)  In  the  "quick  vinegar  pro- 
cess," impure  dilute  alcohol  is  oxidized  to  acetic  acid  by  exposing  it 
to  an  excess  of  air.  The  operation  is  conducted  in  tall  vats  or  casks 
filled  with  beechwood  shavings  soaked  in  old  vinegar.  Holes  at  the 
bottom  and  top  allow  air  to  enter  and  escape  freely.  The  alcoholic 
solution  is  introduced  at  the  top,  trickles  through  the  shavings,  and 
collects  at  the  bottom.  In  its  passage  it  comes  in  contact  with  the 
ferment  and  oxygen,  and  is  partially  converted  into  vinegar.  The 
operation  is  repeated  until  the  change  is  complete.  Thus  prepared, 
the  vinegar  lacks  the  flavor,  odor,  and  color  of  cider  vinegar,  but  these 
deficiencies  may  be  artificially  supplied. 


OTHER  CARBON  COMPOUNDS        319 

Acetates.  —  Acetic  acid  is  a  monobasic  acid,  and  forms  only 
one  series  of  salts  —  the  acetates.  They  are  prepared,  like 
other  salts,  by  the  interaction  of  the  acid  and  carbonates, 
hydroxides,  metals,  etc.  The  metallic  acetates  are  usually 
crystalline  solids,  which  readily  yield  acetic  acid  when 
treated  with  sulphuric  acid.  Most  of  them  contain  water  of 
crystallization,  and  several  are  poisonous. 

Acetates  have  many  applications.  Sodium  acetate,  NaC2H3O2.3  H2O, 
is  a  white  crystalline  solid,  used  in  preparing  pure  aectic  acid  and  in 
the  manufacture  of  dyestuffs.  Lead  acetate,  Pb(C2H3O2)2,  is  a  white 
crystalline  solid,  used  in  dyeing  and  in  making  a  yellow  pigment.  Its 
sweet  taste  led  to  the  common  name  of  "  sugar  of  lead."  Aluminium 
acetate,  A1(C2H3O2)3,  is  not  known  in  the  pure  state,  but  an  impure 
solution,  known  as  "  red  liquor,"  is  extensively  used  in  dyeing  and  in 
calico  printing.  Iron  acetates  are  sold  in  solution  as  a  complex  black 
liquid,  known  as  "  iron  liquor,"  which  is  used  in  dyeing  black  silks 
and  cottons,  and  in  calico  printing.  (See  Mordants.)  A  complex 
copper  acetate,  2  Cu(C2H3O2)2+CuO,  called  verdigris,  is  used  in  making 
blue  paint.  Another  complex  acetate  of  copper  and  arsenic  is  Paris 
green ;  it  is  used  to  kill  potato  bugs  and  other  injurious  insects. 

Other  Organic  Acids  are  oxalic,  lactic,  malic,  butyric,  stearic, 
palmitic,  oleic,  tartaric,  and  citric. 

Oxalic  Acid,  C2H2O4,  occurs  as  a  calcium  salt  (CaC204)  in  many 
plants,  e.g.  rhubarb.  Oxalic  acid  is  a  white  solid,  very  soluble  in 
water,  from  which  it  crystallizes  with  two  molecules  of  water  of 
crystallization  (C2H2O4 .  2  H2O).  It  is  very  poisonous.  The  acid 
and  some  of  its  salts  decompose  iron  rust  and  inks  containing  iron, 
and  are  used  to  remove  stains. 

Lactic  Acid,  C3H6O3,  occurs  in  sour  milk  (see  page  326). 

'Malic  Acid,  C4H6O5,  occurs  free  or  as  salts  in  many  fruits  and 
in  parts  of  vegetables. 

Butyric  Acid,  C4H8O2,  occurs  in  rancid  butter.  Stearic  Acid, 
Ci8H36O2,  and  Palmitic  Acid,  Ci6H32O2,  are  white  solids.  Oleic 
Acid,  Ci8H34O2,  is  an  oily  liquid.  Derivatives  of  these  four  acids 
occur  in  fats  and  oils  (see  page  322). 

Tartaric  Acid,  C4H6O6,  occurs  as  the  acid  potassium  salt 
(HKC4H4O6)  in  grapes  and  other  fruits.  During  the  fermentation 
of  grape  juice,  impure  acid  potassium  tartrate  is  deposited  in  the 
casks.  From  this  argol  or  crude  tartar  the  acid  itself  is  prepared. 


320  INORGANIC   CHEMISTRY 

Tartaric  acid  is  a  white  crystalline  solid,  soluble  in  water  and 
alcohol.  It  is  used  in  dyeing,  and  as  one  ingredient  of  Seidlitz 
powders.  In  these  and  similar  mixtures  it  serves  to  decompose  the 
other  ingredient,  which  is  a  carbonate.  (See  Sodium  Bicarbonate.) 

Tartaric  acid  is  dibasic  and  forms  two  classes  of  salts.  Purified 
acid  potassium  tartrate  obtained  from  argoi  is  commonly  known  as 
cream  of  tartar  (HKC4H4Oe).  It  is  extensively  used  in  the  manu- 
facture of  baking  powders.  These,  as  a  rule,  are  essentially  mix- 
tures of  cream  of  tartar,  sodium  bicarbonate  (HNaCOs),  and  a  little 
starch.  When  moistened  by  dough,  the  baking  powder  dissolves, 
the  acid  salt  and  the  carbonate  interact  and  liberate  carbon  dioxide. 
This  gas  escapes  slowly  through  the  dough,  thereby  puffing  it  up 
and  making  it  porous.  (See  Sodium  Bicarbonate.)  Tartar  emetic  is 
potassium  antimonyl  tartrate  (KSbOC4H4O6).  It  is  used  as  a  medi- 
cine and  to  some  extent  in  dyeing.  Rochelle  salt  is  potassium 
sodium  tartrate  (KNaC4H4O6). 

Citric  Acid,  C6H8O7,  occurs  abundantly  in  lemons  and  oranges, 
and  in  small  quantities  in  currants,  gooseberries,  and  raspberries.  It 
is  a  white  crystalline  solid,  very  soluble  in  water.  The  taste  is  sour, 
but  pleasant.  The  acid  and  its  magnesium  salt  are  used  as  medi- 
cines. The  acid  itself  is  used  in  calico  printing.  Citric  acid  is  tribasic. 

ESTERS 

Esters  are  compounds  of  carbon,  hydrogen,  and  oxygen 
closely  related  to  alcohols  and  organic  acids.  Thus,  when 
ethyl  alcohol,  acetic  acid,  and  concentrated  sulphuric  acid 
are  mixed  and  warmed,  ethyl  acetate  is  one  product.  The 
essential  change  is  represented  thus :  — 

C2H5 .  OH  +  CH3 .  COOH  =  CH3 .  COOC2H5  +  H2O 

Ethyl  Alcohol  Acetic  Acid  Ethyl  Acetate        Water 

Ethyl  acetate  has  a  pleasant,  fruitlike  odor,  and  its  forma- 
tion in  this  way  is  a  simple  test  for  alcohol  or  acetic  acid. 
Ethyl  acetate  is  analogous  to  sodium  acetate,  i.e.  the  organic 
salt  contains  the  radical  ethyl,  while  the  metallic  salt  contains 
sodium.  The  fatty  acids,  as  well  as  those  of  other  series, 
form  many  esters  of  special  interest.  Some  occur  naturally 


OTHER  CARBON  COMPOUNDS  321 

in  fruits  and  flowers,  and  in  many  cases  give  the  fragrance 
and  flavor.  Others  are  prepared  artificially  and  used  as  the 
characteristic  ingredient  of  cheap  flavoring  extracts,  per- 
fumery, and  beverages.  Ethyl  butyrate  has  the  taste  and 
fragrance  of  pineapples,  amyl  acetate  of  bananas,  amyl 
valerate  of  apples,  methyl  salicylate  of  wintergreen. 

FATS,  GLYCERIN,  AND  SOAP 

General  Relations.  —  Natural  fats  and  oils  are  essentially 
mixtures  of  stearin,  palmitin,  and  olein.  Beef  and  mutton 
fat  are  chiefly  stearin,  lard  is  mainly  palmitin  and  olein,  while 
oils  such  as  olive  oil  are  largely  olein.  Stearin  and  palmitin 
are  solids  at  the  ordinary  temperature,  but  olein  is  a  liquid. 
These  three  compounds  —  stearin,  palmitin,  and  olein  — 
are  esters  of  their  corresponding  acids  and  the  alcohol  gly- 
cerin. They  are  analogous  to  ethyl  acetate.  The  radical 
of  glycerin  is  glyceryl,  C3H5.  Stearin  is  glyceryl  stearate, 
palmitin  is  glyceryl  palmitate,  and  olein  is  glyceryl  oleate. 
Natural  fats  and  oils,  therefore,  are  mixtures  of  these  and 
similar  esters.  Glycerin  is  a  triacid  alcohol  containing  three 
hydroxyl  (OH)  groups.  Like  ordinary  alcohol,  it  interacts 
with  the  fatty  acids  and  forms  esters.  The  latter,  as  we  have 
just  seen,  are  the  fats.  Now,  when  fats  are  heated  with  very 
hot  steam  or  with  sulphuric  acid,  they  are  changed  into 
glycerin  and  the  corresponding  acids.  Thus,  with  stearin 
the  change  is  — 

(Ci7H35 .  COO)3C3H5  +  3  H20  =  C8HB(OH)j  +  3  Ci7H35 .  COOH 

Stearin  Glycerin  Stearic  Acid 

But  if  fats  are  boiled  with  sodium  hydroxide  or  a  similar 
alkali,  glycerin  and  an  alkaline  salt  of  the  corresponding 
acid  are  formed.  Soap  is  a  mixture  of  such  alkaline  salts. 
In  a  few  words,  the  general  relations  are  these :  (1)  fats  are 
esters;  (2)  treated  with  steam  or  acid,  fats  form  glycerin 


322  INORGANIC  CHEMISTRY 

and  fatty  acids;   (3)  treated  with  alkalies,  fats  form  glycer- 
in and  soap. 

Natural  Fats  and  Oils  are  often  complex  mixtures.  The 
solid  fats,  as  already  stated,  are  rich  in  stearin  and  palmitin. 
Tallow  is  chiefly  stearin,  but  human  fat  and  palm  oil  are 
largely  palmitin.  The  soft  and  liquid  fats  and  oils  contain 
considerable  olein,  as  a  rule.  The  proportion  of  olein  deter- 
mines the  consistency  of  the  fats  and  oils.  Thus,  olive  oil 
contains  72  per  cent  of  olein  (and  a  similar  fat)  and  about 
28  per  cent  of  stearin  and  palmitin.  The  specific  character 
of  many  fats  and  oils  is  due  mainly  to  a  small  proportion  of 
certain  fats.  These  fats  correspond  to  uncommon  acids  in 
the  fatty,  oleic,  and  other  series.  Butter,  for  example,  con- 
sists of  the  fats  corresponding  to  the  following  acids :  palmitic, 
stearic,  oleic,  butyric,  capric,  and  caproic.  The  last  three, 
together  with  traces  of  other  substances,  give  butter  its 
pleasant  flavor.  Oleomargarine  and  other  substitutes  for 
butter  resemble  real  butter  very  closely  in  composition. 
Artificial  butter,  however,  lacks  the  flavor  of  the  real  butter, 
but  it  is  "  probably  just  as  nutritious,  although  perhaps 
not  quite  so  easily  digested." 

Glycerin,  C3H803  or  C3H5 .  (OH)3,  is  a  thick,  sweet  liquid. 
It  mixes  readily  with  water  and  with  alcohol  in  all  propor- 
tions, and  absorbs  moisture  from  the  air.  Heated  in  the  air 
it  decomposes  and  gives  off  irritating  gases,  like  those  pro- 
duced by  burning  fat. 

Glycerin  is  used  to  make  nitroglycerin  (see  below),  toilet 
soaps,  and  printer's  ink  rolls ;  it  is  also  used  as  a  solvent, 
a  lubricator,  a  preservative  for  tobacco,  and  certain  foods, 
a  sweetening  substance  in  certain  liquors,  preserves,  and 
candy ;  as  a  cosmetic ;  and  owing  to  its  non-volatile  and 
non-drying  properties,  it  is  used  as  an  ingredient  of  certain 
inks  and  oils. 


OTHER  CARBON  COMPOUNDS  323 

Glycerin  is  a  by-product  in  the  manufacture  of  soap,  or  it  is 
made  directly  by  decomposing  fats  with  steam  under  pressure  or 
with  lime.  All  these  methods  involve  the  chemical  change  described 
above,  viz.  the  decomposition  of  an  ester  (the  fat)  into  the  cor- 
responding alcohol  (glycerin)  and  a  mixture  of  fatty  acids.  By 
skillful  treatment  the  glycerin  is  freed  from  the  water  and  im- 
purities. 

As  already  stated,  glycerin  is  an  alcohol,  and  for  this  reason  it 
is  often  called  glycerol.  When  treated  with  a  mixture  of  concen- 
trated nitric  and  sulphuric  acids,  it  forms  an  ester  commonly  known 
as  nitroglycerin  (C3H5(ONO2)3).  This  is  a  yellow,  heavy,  oily 
liquid.  It  is  the  well-known  explosive,  and  is  also  an  ingredient  of 
some  other  explosives.  When  kindled  by  a  flame,  it  burns  with- 
out explosion  ;  but  if  it  is  compressed,  detonated,  or  heated  to 
about  250°  C.,  it  explodes  violently.  Nitroglycerin  is  dangerous  to 
handle  and  transport,  and  is  usually  mixed  with  some  porous  sub- 
stance, such  as  infusorial  earth.  In  this  form  it  is  called  dynamite. 
Other  explosives  contain  nitroglycerin,  e.g.  blasting  gelatin  and 
cordite. 

Soap,  as  already  stated,  is  a  mixture  of  alkaline  salts  of 
organic  acids,  mainly  stearic  and  palmitic  acids.  Soap  is 
made  by  boiling  fats  with  sodium  hydroxide  or  potassium 
hydroxide.  This  process  is  called  saponification.  Sodium 
hydroxide  produces  hard  soap,  consisting  chiefly  of  sodium 
palmitate,  sodium  stearate,  and  sodium  oleate.  Potassium 
hydroxide  produces  a  soft,  semi-fluid  soap,  which  contains 
mainly  the  corresponding  potassium  salts.  The  chemical 
change,  as  already  stated,  consists  in  the  transformation  of 
an  ester  (fat)  into  glycerin  and  an  alkaline  salt.  In  the  case 
of  pure  stearin  (glyceryl  stearate)  the  change  may  be  repre- 
sented thus  :  — 


Stearin  Sodium  Sodium  Glycerin 

Hydroxide  Stearate 

The  fats  used  in  soap  making  vary.     Tallow,  lard,  palm  oil, 
and  cocoanut  oil  make  white  soaps.     Grease,  together  with 


324  INORGANIC  CHEMISTRY 

tallow,  palm  oil,  cottonseed  oil,  and  rosin,  make  yellow  soaps. 
Olive  oil  is  used  for  making  castile  soap. 

Most  soaps  are  manufactured  by  the  boiling  process.  The  fat 
and  alkali  are  boiled  in  a  huge  kettle.  This  operation  produces  a 
thick,  frothy  mixture  of  soap,  glycerin,  and  alkali.  At  the  proper 
time  salt  is  added,  thereby  causing  the  soap  to  separate  and  rise  to 
the  top.  The  liquid  beneath  is  drawn  off,  and  from  it  glycerin  is 
extracted.  The  soap  is  often  boiled  again  with  rosin  or  cocoanut  oil ; 
then  mixed,  if  desired,  with  perfume,  coloring  matter,  or  some 
filling  material  (such  as  sodium  silicate,  sand,  or  borax).  Floating 
soaps  are  made  by  forcing  air  into  the  semi-solid  soap  before  cool- 
ing. The  best  soaps  do  not  contain  unchanged  fat  or  "  free  alkali," 
i.e.  sodium  hydroxide. 

The  cleansing  action  of  soap  is  probably  due  to  two  causes. 
(1)  Soap  hydrolyzes  with  water  and  the  liberated  sodium  hydroxide 
acts  upon  the  grease  and  oil  that  are  mixed  with  the  dirt.  (2)  Soap 
causes  oils  to  form  an  emulsion  which  is  readily  removed  by  water. 
Doubtless  the  second  cause  is  the  more  efficient. 

CARBOHYDRATES 

Sugar.  —  The  popular  term  sugar  means  almost  any 
sweet  substance  found  in  fruits,  nuts,  vegetables,  sap  of 
trees,  etc.,  though  it  is  usually  restricted  to  the  ordinary 
white  sugar  obtained  from  sugar  cane  and  sugar  beet. 
Chemically,  there  are. many  sugars,  each  having  a  definite 
constitution.  The  most  important  is  ordinary  sugar,  which 
is  also  called  cane  sugar,  sucrose,  and  saccharose.  Other 
important  sugars  are  dextrose,  levulose,  lactose,  and  maltose. 

Cane  Sugar,  C^H^On,  is  widely  distributed  in  nature, 
being  found  in  the  sugar  cane,  sugar  beet,  sugar  maple, 
Indian  corn,  sorghum,  most  sweet  fruits,  many  nuts,  blos- 
soms of  flowers,  and  honey.  The  main  source  of  cane  sugar 
is  the  sugar  cane  and  sugar  beet. 

Cane  sugar  is  a  white,  crystalline  solid.  It  is  very  soluble 
in  water,  one  part  of  water  dissolving  about  three  times  its 


OTHER  CARBON  COMPOUNDS  325 

weight  of  sugar  at  ordinary  temperatures.  If  heated  to  about 
160°  C.,  it  melts,  and  on  cooling  becomes  a  glassy  solid.  As 
the  temperature  is  raised,  the  solid  begins  to  decompose, 
and  at  about  210°  C.  water  is  given  off  and  a  brown  substance 
called  caramel  is  formed.  Further  heating  produces  a  black 
porous  mass  of  carbon  called  sugar  charcoal. 

The  manufacture  of  sugar  from  sugar  cane  and  sugar  beets 
involves  two  main  operations.  (1)  In  the  preparation  of  raw  sugar 
from  sugar  cane  the  juice  obtained  by  crushing  the  cane  is  first  boiled 
with  a  weak  calcium  hydroxide  solution  to  neutralize  acids,  remove 
impurities,  and  prevent  fermentation,  next  freed  from  excess  of  lime 
by  carbon  dioxide,  and  finally  filtered  through  bone  black.  The 
purified  juice  is  then  evaporated  in  vacuum  pans  until  the  sugar  be- 
gins to  crystallize  from  the  cooled  liquid.  The  crystals  are  then 
separated  from  the  brown  liquid  by  a  centrifugal  machine.  The 
liquid  is  the  familiar  molasses.  In  the  preparation  of  raw  sugar 
from  sugar  beets  the  washed  beets  are  cut  into  slices  and  soaked  in 
water  to  dissolve  the  sugar.  The  solution  is  treated  much  like  cane 
sugar  solutions.  (2)  Raw  sugar  is  dark  colored,  and  must  be  re- 
fined before  it  is  suitable  for  most  uses.  The  raw  sugar  is  first  dis- 
solved in  water,  and  lime  and  other  substances  are  added  to  gather 
the  impurities  into  a  scum  or  clot.  The  colored  liquid  is  next 
filtered,  first  through  cloth  bags  and  then  through  bone  black. 
The  filtered  sirup  is  evaporated  in  large  vacuum  pans  until  a  sample 
deposits  the  right  size  crystals.  The  crystals  of  sugar  are  separated 
from  the  sirup  by  centrifugal  machines,  then  dried  and  separated 
in  a  heated  tube  called  a  granulator.  Hence  the  name  granulated 
sugar. 

Dextrose  and  Levulose.  —  When  sucrose  is  heated  with 
dilute  acids,  the  two  sugars  dextrose  and  levulose  are  formed. 
The  chemical  change  is  an  example  of  hydrolysis  and  may 
be  represented  thus  :  — 

CisHaOu     +      H20     =      C6H1206     +      C6H1206 

Sucrose  Dextrose  Levulose 

The  same  change  is  brought  about  by  an  enzyme  called 
invertase.  Dextrose  is  a  white  solid  about  three  fifths  as 


326  INORGANIC  CHEMISTRY 

sweet  as  sucrose.  Dextrose  is  found  in  honey  and  in  many 
fruits,  especially  grapes,  and  is  sometimes  called  grape  sugar. 
Another  name  for  it  is  glucose.  Levulose  is  also  a  sweet, 
white  solid  found  in  fruits  and  honey,  and  is  often  associated 
with  dextrose.  It  is  sometimes  called  fructose  or  fruit  sugar. 

Commercial  glucose  contains  about  40  to  50  per  cent  of  dextrose. 
It  is  manufactured  by  heating  starch  with  dilute  sulphuric  acid. 
The  starch  is  first  changed  into  a  sweet  solid  called  dextrin,  then  into 
dextrose,  and  if  the  process  is  carried  far  enough,  the  product  is  a 
hard,  waxlike  solid  known  as  commercial  grape  sugar,  which  is 
almost  pure  dextrose.  Glucose-  is  an  inexpensive  substitute  for 
sucrose  and  is  extensively  used  in  making  candy,  jellies,  sirups,  and 
other  sweet  mixtures. 

Dextrose,  and  also  levulose,  is  converted  by  yeast  into  ethyl 
alcohol  and  carbon  dioxide  (see  Alcohol).  Dextrose  and  levulose 
are  reducing  agents.  An  alkaline  solution  of  dextrose  is  used  to 
reduce  a  silver  solution  and  deposit  the  silver  as  a  bright  film  in 
making  reflectors,  mirrors,  Dewar  flasks,  and  thermos  bottles. 
It  also  reduces  a  strongly  alkaline  mixture  of  copper  sulphate  and 
sodium  potassium  tartrate,  known  as  Fehling's  solution.  When  this 
solution  is  boiled  with  dextrose  (or  any  other  reducing  sugar),  a 
reddish  copper  compound  (cuprous  oxide,  Cu20)  is  formed.  This 
experiment  is  often  used  as  a  test  for  dextrose  and  similar  sugars. 
Solutions  of  dextrose  and  levulose  rotate  the  plane  of  polarized 
light  —  dextrose  to  the  right  and  levulose  to  the  left.  That  is, 
when  their  solutions  are  placed  in  a  sugar-polariscope  and  examined, 
the  light  instead  of  passing  entirely  through  the  instrument  is 
extinguished ;  and  in  order  to  bring  about  illumination  again,  the 
plane  of  the  polarized  light  must  be  rotated  a  certain  number  of 
degrees  in  order  to  compensate  for  the  rotation  caused  by  the  sugar 
solution.  By  means  of  this  instrument  valuable  information  can 
be  obtained  about  the  kind  and  proportion  of  sugar  in  solutions. 

Lactose  (milk  sugar,  Ci2H22On  .  H2O)  occurs  in  the  milk  of 
mammals.  Cow's  milk  contains  on  the  average  4.88  per  cent  of 
lactose.  Lactose  is  not  so  sweet  or  soluble  as  cane  sugar.  When 
milk  sours,  its  lactose  changes  into  alcohol  and  lactic  acid.  The 
acid  causes  the  sour  taste  and  also  assists  in  curdling  the  milk,  i.e. 
in  changing  the  casein  into  a  clot  or  curd.  Casein  is  used  in  making 
cheese. 


OTHER  CARBON  COMPOUNDS  327 

• 

Maltose  is  formed  from  starch  by  malt,  hence  the  name  maltose. 
The  transformation  is  caused  by  the  enzyme  diastase.  Maltose 
ferments  readily  with  yeast,  forming  alcohol  and  carbon  dioxide, 
and  is  manufactured  in  large  quantities  for  the  commercial  produc- 
tion of  alcohol  and  fermented  liquors.  With  dilute  acids,  maltose 
forms  dextrose  by  hydrolysis.  Like  lactose,  maltose  is  a  sweet 
solid,  very  soluble  in  water,  from  which  it  forms  crystals  (Ci2H22Oii . 
HzO) ;  its  solution  turns  the  plane  of  polarized  light  to  the  right  and 
reduces  Fehling's  solution. 

Starch  is  widely  distributed  in  the  vegetable  kingdom. 
It  is  found  in  wheat,  corn,  and  all  other  grains  ;  in  potatoes, 
beans,  peas,  and  similar  vegetables ;  and  in  large  quantities 
in  rice,  sago,  tapioca,  and  nuts.  Many  parts  of  plants  con- 
tain starch. 

Starch  is  a  white  mass,  as  usually  seen.  But  under  the 
microscope  it  is  found  to  consist  of  oval  grains  varying  some- 


FIG.  50.  —  Starch  grains  (magnified)  —  wheat  (left) ,  rice  (center) ,  corn  (right). 

what  with  the  source  (Fig.  50) .  Starch  is  only  very  slightly 
soluble  in  water.  But  if  heated  with  water,  the  grains  swell 
and  burst,  partially  dissolve,  and  form  a  solution  which,  when 
cold,  becomes  the  familiar  starch  paste.  Starch  in  solution 
is  turned  blue  by  iodine,  and  its  presence  in  many  vegetables 
and  foods  can  be  readily  shown  by  grinding  the  substance  in 
a  mortar  with  cold  water  and  adding  a  drop  of  dilute  iodine 
solution.  The  composition  of  starch  corresponds  to  the 
formula  (C8Hi0O5)n. 


328  INORGANIC  CHEMISTRY 

• 

Bread.  —  Wheat  flour  contains  about  70  per  cent  of  starch. 
The  remainder  is  chiefly  water  and  gluten,  though  small 
quantities  of  mineral  matter  and  fat  are  present.  In  making 
bread,  the  flour,  water,  and  yeast  are  thoroughly  mixed  into 
dough,  which  is  put  in  a  warm  place  to  rise.  Fermentation 
begins  at  once.  The  enzymes  change  the  starch  into  dextrose 
or  a  similar  fermentable  substance  which  undergoes  fer- 
mentation, forming  alcohol  and  carbon  dioxide.  The  gases 
escape  in  part  through  the  dough,  which  becomes  light  and 
porous.  When  the  dough  is  baked,  the  heat  kills  the  yeast 
plant  and  fermentation  stops ;  but  the  alcohol,  carbon  diox- 
ide, and  some  water  escape  and  puff  up  the  mass  still  more. 
The  heat,  however,  soon  hardens  the  starch,  gluten,  etc., 
into  a  firm  but  porous  loaf. 

Cellulose,  (C6H10O5)ra,  is  the  basis  of  the  cells  of  plants. 
Wood,  cotton,  linen,  and  paper  are  largely  cellulose.  Pure 
cellulose  is  a  white  substance,  insoluble  in  most  liquids,  but 
soluble  in  a  mixture  of  ammonia  and  copper  oxide.  Concen- 
trated sulphuric  acid  dissolves  it  slowly ;  and  if  the  solution 
is  diluted  and  boiled,  the  cellulose  is  changed  into  a  mixture 
of  glucose  and  dextrin.  Sulphuric  acid  of  a  special  strength, 
if  quickly  and  properly  applied  to  paper,  changes  it  into  a 
tougher  form  called  parchment  paper.  Ordinary  paper 
consists  chiefly  of  cellulose  matted  together. 

Cellulose  with  nitric  acid  forms  cellulose  nitrates.  One  of  the 
cellulose  nitrates  is  gun  cotton.  It  looks  like  ordinary  cotton,  and 
may  be  spun,  woven,  and  pressed  into  cakes.  It  burns  quickly, 
if  unconfined  ;  but  when  ignited  by  a  percussion  cap  or  when  burned 
in  a  confined  space,  gun  cotton  explodes  violently.  It  is  used  in 
blasting  and  for  torpedoes  and  submarine  mines.  A  mixture  of 
gun  cotton,  alcohol,  and  ether  forms  a  transparent  solid  called  smoke- 
less powder ;  when  exploded  it  yields  colorless  gases.  A  solution 
of  certain  cellulose  nitrates  in  a  mixture  of  alcohol  and  ether  is 
called  collodion.  It  is  used  in  preparing  certain  photographic 


OTHER  CARBON  COMPOUNDS  329 

material  and  as  a  coating  for  wounds.     A  mixture  of  camphor  and 
cellulose  nitrates  is  called  celluloid. 

BENZENE  AND  ITS  DERIVATES 

Benzene,  C6H6,  is  a  colorless  liquid,  lighter  than  water 
(sp.  gr.  .88  at  20°  C.)  and  boils  at  80°  C.  It  burns  with  a 
luminous,  smoky  flame,  owing  to  its  richness  in  carbon. 
Ordinary  illuminating  gas  owes  its  luminosity  partly  to  ben- 
zene. It  dissolves  fats,  resins,  iodine,  sulphur,  and  rubber. 
Benzene  is  sometimes  called  benzol.  It  should  not  be  con- 
fused with  benzine,  which  is  a  mixture  of  hydrocarbons  de- 
rived from  petroleum.  Benzene  is  chiefly  used  in  preparing 
its  derivatives. 

Nitrobenzene,  C6H5 .  N02,  is  a  yellow  liquid  formed  by  the  inter- 
action of  benzene  and  nitric  acid.  The  equation  for  the  chemical 
change  is :  — 

C6H6     +     HN03    -    C6H5.N02    +     H2O 

Benzene  Nitric  Nitro-  Water 

Acid  benzene 

It  is  volatile,  and  the  vapor,  which  is  poisonous,  has  the  odor  of 
bitter  almonds.     It  is  chiefly  used  in  the  manufacture  of  aniline. 

Aniline,  C6H5  .  NH2,  is  an  oily  liquid,  slightly  heavier  than  water. 
It  is  prepared  on  a  large  scale  by  reducing  nitrobenzene  with  nascent 
hydrogen,  thus :  — 

CeH6 .  NO2     +     6  H     =     C6H5 .  NH2     +     2  H2O 

Nitrobenzene  Hydrogen  Aniline  Water 

From  aniline  are  made  many  compounds  known  as  aniline  dyes. 

Phenol,  C6H5  .  OH,  is  a  white  crystalline  solid.  It  has  a  smoky 
odor,  is  poisonous,  and  burns  the  skin.  Coal  tar  is  the  source  of 
phenol.  A  solution  of  phenol  in  water,  popularly  called  carbolic 
acid,  is  used  as  a  disinfectant. 

Naphthalene,  Ci0H8,  is  a  white,  lustrous,  crystalline  solid  ob- 
tained from  coal  tar.  It  has  a  penetrating,  unpleasant  odor,  and  is 
used  as  a  substitute  for  camphor  under  the  name  of  "  moth  balls." 
Large  quantities  of  naphthalene  are  used  in  making  indigo. 


330  INORGANIC  CHEMISTRY 

Anthracene,  Ci4H10,  is  a  white  crystalline  solid,  and  is  obtained 
from  coal  tar.  It  is  one  of  the  most  important  hydrocarbons,  be- 
cause from  it  alizarin  is  made.  Alizarin  is  a  valuable  dyestuff, 
because  it  produces  brilliant,  fast  colors  with  different  mordants. 

CYANOGEN  AND  ITS  DERIVATIVES 

Cyanogen,  (CN)2,  is  a  colorless  gas,  has  the  odor  of  peach 
kernels,  is  poisonous,  and  burns  with  a  purplish  flame. 

Hydrocyanic  or  prussic  acid,  HCN,  is  prepared  by  heating 
a  cyanide  with  sulphuric  acid.  The  acid  smells  like  peach 
kernels,  and  is  extremely  poisonous.  It  is  a  feeble  acid  and 
dissociates  slightly  into  H+  and  CN~. 

Potassium  cyanide,  KCN,  and  sodium  cyanide,  NaCN, 
are  white,  deliquescent  solids.  They  are  very  poisonous. 
Large  quantities  are  used  in  gold  and  silver  plating  and  in 
the  cyanide  process  of  extracting  gold  from  its  ores,  as  de- 
scribed under  that  metal.  The  alkaline  reaction  exhibited 
by  a  solution  of  potassium  cyanide  is  due  to  hydrolysis. 

Potassium  sulphocyanate  or  thiocyanate,  KCNS,  is  a  white 
crystalline  salt,  which  produces  a  deep  red  solution  (due  to 
ferric  sulphocyanate)  with  solutions  containing  ferric  ion. 
(See  Ferric  Compounds.) 

PROTEINS 

Proteins  are  complex  nitrogenous  compounds  often  desig- 
nated by  the  term  .protein  (formerly  proteid).  Besides  nitro- 
gen, they  contain  carbon,  hydrogen,  and  oxygen,  usually 
sulphur,  often  phosphorus,  and  occasionally  iron.  Common 
proteins  contain  15  to  18  per  cent  of  nitrogen.  They  burn 
with  a  disagreeable  odor  and  liberate  ammonia.  When 
proteins  putrefy,  poisonous  substances  called  ptomaines  are 
often  produced. 

Proteins  constitute  the  principal  part  of  the  tissue  of  the 


OTHER  CARBON  COMPOUNDS  331 

cells  of  our  bodies.  The  body  of  an  average  man  is  about  18 
per  cent  protein.  Protein  in  some  form  must  be  a  part  of 
the  food  of  animals,  since  protein  serves  to  replace  worn- 
out  tissue  and  contribute  new  tissue  for  the  growth,  of  the 
body.  Important  groups  of  proteins  are  albumins,  globulins, 
glutelins  and  protamins,  albuminoids,  phosphoproteins,  and 
hemoglobins.  Members  of  these  groups  are  widely  distrib- 
uted, being  found  in  eggs,  milk,  seeds  of  plants  (e.g.  beans 
and  peas),  cereals  (e.g.  wheat,  rye,  and  barley),  muscle, 
blood,  and  animal  tissues. 

PROBLEMS 

1.  Calculate  the  affinity  constant  of  a  molar  solution  of  acetic* 
acid  in  which  the  proportion  ionized  is  .0041.  Ans.  .0000169. 

2.  What  weight  of  potassium  hydroxide  will  neutralize  (a)  100 
gm.  of  acetic  acid  and  (6)  10  gm.  of  oxalic  acid  (dibasic)  ? 


CHAPTER  XVIII 
Sulphur  and  its  Compounds 

SULPHUR  has  been  known  for  ages.  The  alchemists  re- 
garded it  as  one  of  the  primary  forms  of  matter.  The  ele- 
ment and  its  compounds  have  always  played  an  important 
part  in  the  development  of  many  industries. 

Occurrence.  —  Sulphur,  free  and  combined,  is  abundant 
and  widely  distributed.  Free  or  native  sulphur  is  found 
usually  in  volcanic  regions.  There  are  also  beds  associated 
with  gypsum  (calcium  sulphate).  It  is  believed  that  some 
deposits  were  formed  by  the  reduction  of  gypsum  by  micro- 
organisms. 

Several  important  metallic  ores  are  native  sulphides,  e.g. 
lead  sulphide  (PbS),  zinc  sulphide  (ZnS),  and  those  of  mer- 
cury, antimony,  and  copper.  Iron  sulphide  (FeS2)  is  also 
plentiful.  The  most  abundant  sulphates  are  calcium  sul- 
phate (CaSO4),  barium  sulphate  (BaSO4),  and  magnesium 
sulphate  (MgSO4).  Volcanic  gases  often  contain  sulphur 
dioxide  (SO2)  and  hydrogen  sulphide  (H2S).  The  latter  is 
also  found  in  sulphur  springs.  Sulphur  is  a  constituent  of 
protein,  and  hence  is  present  in  many  kinds  of  animal  and 
vegetable  matter,  e.g.  eggs  and  mustard.  Some  varieties  of 
petroleum  and  coal  contain  sulphur  compounds. 

Source.  —  Until  about  1903  Sicily  furnished  most  of  the 
sulphur,  but  the  larger  part  of  the  world's  supply  now  comes 
from  Japan  and  the  United  States  (especially  Louisiana). 
Some  of  the  sulphur  of  commerce  is  obtained  by  roasting 

332 


SULPHUR  AND  ITS  COMPOUNDS 


333 


iron  pyrites  (FeS2)  as  in  the  manufacture  of  sulphuric  acid. 
Small  amounts  are  recovered  from  the  calcium  sulphide 
waste  of  the  Leblanc  soda  process  (see  Sodium  Carbonate), 
and  from  the  residues  of  the  iron  oxide  used  to  purify  illu- 
minating gas. 

Extraction.  —  For  many  years 
sulphur  has  been  extracted  from 
the  impure  native  sulphur  in  Sicily 
by  a  primitive  process.  The  sul- 
phur ore  is  brought  to  the  surface 
by  laborers,  piled  loosely  in  a  heap, 
and  covered  with  powdered  or 
burnt  ore  or  with  earth.  The 
heap  is  ignited  at  the  bottom,  and 
the  heat  produced  by  the  combus- 
tion of  some  of  the  sulphur  melts 
the  rest,  which  runs  out  at  the 
bottom.  In  Louisiana,  which  now 
furnishes  most  of  the  sulphur  used 


FIG.  51.  —  Section  of  a  set  of 
pipes  for  winning  sulphur. 
The  water  enters  through  the 
spaces  A  and  B  and  melts  the 
sulphur,  which  rises  part  way 
in  C  and  is  forced  to  the  sur- 
face by  compressed  air  intro- 
duced through  the  innermost 
pipe. 


in  the  United  States,  the  deposits 
are  about  half  a  mile  in  diameter  and  500  feet  thick  and 
are  located  about  800  feet  beneath  the  surface.  They  are 
reached  by  wells  consisting  essentially  of  a  set  of  four  con- 
centric pipes  (Fig.  51)  sunk  through  a  mass  of  clay,  quick- 
sand, and  rock.  Water  heated  (under  pressure)  to  about 
170°  C.  is  pumped  down  the  two  outer  pipes  (A,  B).  After 
some  of  the  sulphur  is  melted,  compressed  air  is  forced  down 
through  the  innermost  pipe  into  the  molten  mass,  thereby 
forcing  the  mixture  of  air  and  melted  sulphur  up  through 
the  pipe  C.  The  liquid  flows  into  huge  wooden  bins  where 
it  solidifies.  The  wells  are  very  powerful,  a  single  well 
often  pumping  500  tons  of  sulphur  daily.  The  annual 
production  is  over  200,000  tons. 


334 


INORGANIC  CHEMISTRY 


Purification.  —  Louisiana  sulphur  is  about  99  per  cent 
pure.  Crude  Sicilian  sulphur  requires  purification.  This 
is  accomplished  by  the  apparatus  shown  in  Figure  52.  The 
crude  sulphur  is  melted  in  B,  and  flows  into  the  iron  cylin- 


FIG.  52.  —  Apparatus  for  purifying  sulphur. 

der  A.  Here  it  is  heated,  and  the  vapors  pass  into  the 
large  brick  chamber,  provided  with  a  tap  C,  from  which 
the  liquid  sulphur  may  be  withdrawn.  If  the  distillation 
is  conducted  slowly,  the  sulphur  vapor  condenses  upon  the 
cold  walls  of  the  chamber  as  a  fine  powder,  called  flowers 
of  sulphur.  As  the  operation  continues  the  walls  become 
hot,  and  the  sulphur  collects  on  the  floor  as  a  liquid  which  is 
drawn  off  into  cylindrical  wooden  molds,  forming  roll  sul- 
phur or  brimstone. 

Properties.  —  Ordinary  sulphur  is  a  yellow,  brittle,  crystal- 
line solid.     It  is  practically  insoluble  in  water,  but  most 


SULPHUR  AND  ITS  COMPOUNDS  335 

varieties  dissolve  in  carbon  disulphide,  and  to  some  extent  in 
turpentine,  chloroform,  and  benzene  (C6H6).  Sulphur  does 
not  conduct  heat  well,  the  warmth  of  the  hand  even  causing 
it  to  crackle  and  break  from  the  unequal  expansion.  The 
specific  gravity  of  the  solid  is  about  two.  The  density  of 
the  vapor  varies  with  the  temperature.  At  the  lowest  tem- 
perature at  which  sulphur  can  be  vaporized,  one  molecule 
contains  eight  atoms  (S8),  while  at  800°  C.  it  contains  two 
atoms  (S2).  Dissolved  sulphur  has  the  formula  S8. 

Heated  to  114.5°  C.  sulphur  melts  to  a  thin  pale  yellow 
liquid.  At  about  160°  C.  the  liquid  becomes  dark  brown 
and  viscous,  and  at  about  230°  C.  it  is  black  and  too  thick 
to  be  poured  from  the  vessel.  Heated  still  higher,  the  color 
remains  black  but  the  mass  becomes  thin,  and  finally  at 
about  445°  C.  the  liquid  boils  and  turns  into  yellow  sulphur 
vapor.  When  cooled  slowly,  the  sulphur  undergoes  the  same 
series  of  changes  in  the  reverse  order.  Sulphur  ignites 
readily  and  burns  with  a  pale  blue  flame,  forming  sulphur 
dioxide  gas  (SO2) ;  if  burned  in  oxygen,  a  little  sulphur 
trioxide  (SO3)  is  also  formed.  Finely  divided  sulphur  oxidizes 
in  moist  air,  forming  sulphuric  acid  (H2SO4).  It  combines 
directly  and  readily  with  hydrogen,  carbon,  and  chlorine, 
forming  hydrogen  sulphide,  carbon  disulphide,  and  sulphur 
chlorides  (S2C12  and  SC14).  The  reaction  between  sul- 
phur and  metals  is  often  attended  by  vivid  combustion, 
though  heat  is  necessary  to  start  the  chemical  action.  Thus, 
when  a  mixture  of  flowers  of  sulphur  and  powdered  iron  is 
heated,  the  mass  begins  to  glow  and  soon  becomes  red-hot, 
the  glow  often  spreading  through  the  mass  after  removal 
from  the  flame.  The  product  is  iron  sulphide,  and  the  chemi- 
cal change  is  represented  thus  :  — 

Fe     +     S  FeS 

Iron  Sulphur  Iron 

Sulphide 


336  INORGANIC  CHEMISTRY 

Heated  copper  glows  when  dropped  into  melted  sulphur, 
while  zinc  dust  and  flowers  of  sulphur  combine^  violently. 

Different  Forms  of  Sulphur.  —  Sulphur  exists  in  several 
different  forms,  which  are  crystallized  or  amorphous.  These 
modifications  differ  in  specific  gravity,  solubility,  and  other 
properties.  The  crystallized  forms  belong  to  the  ortho- 
rhombic  and  monoclinic  systems.  (See  Appendix,  §  3.) 
Orthorhombic  sulphur  is  the  form  obtained  by  crystalliza- 
tion from  a  solution  of  carbon  disulphide.  Crystallized 
native  sulphur  and  ordinary  roll  sulphur  are  orthorhombic, 
though  the  latter  usually  consists  of  such  a  mass  of  inter- 
laced crystals  that  the  form  is  obscured.  The  monoclinic 
sulphur  is  the  form  obtained  by  slowly  cooling  molten  sul- 
phur. By  melting  sulphur  in  a  crucible  and  pouring  off  the 
excess  of  liquid  as  soon  as  crystals  shoot  out  from  the  walls 
near  the  surface,  the  interior  of  the  crucible  is  found  to  be 
lined  with  long,  dark  yellow,  shining  needles.  They  are 
monoclinic  crystals  of  sulphur.  After  a  few  days  they  gradu- 
ally become  opaque  and  slowly  change  into  minute  ortho- 
rhombic  crystals.  This  change  is  due  to  the  fact  that 
monoclinic  sulphur  is  in  the  stable  form  only  between  96° 
C.  and  119°  C.  (its  melting  point),  while  orthorhombic  sul- 
phur is  in  the  stable  form  only  below  96°  C.  Above  96°  C. 
orthorhombic  slowly  becomes  monoclinic,  below  96°  C.  mono- 
clinic  slowly  becomes  orthorhombic.  This  temperature  at 
which  the  two  forms  of  sulphur  pass  into  one  another  is 
called  the  transition  point.  These  two  varieties  of  crystal- 
lized sulphur  have  different  properties.  Orthorhombic  sul- 
phur has  the  specific  gravity  2.06  and  melts  at  114.5°  C. 
(if  heated  rapidly) ;  the  corresponding  values  of  monoclinic 
sulphur  are  1.96  and  119.25°  C.  Amorphous  sulphur  is 
formed  by  pouring  boiling  sulphur  into  cold  water.  It  is  a 
tough,  plastic,  rubberlike,  amber-colored  mass,  mostly  in- 


SULPHUR  AND  ITS  COMPOUNDS  337 

soluble  in  carbon  disulphide.  Its  formation  is  due  to  the 
sudden  cooling  of  the  viscous  liquid  sulphur.  It  soon  be- 
comes hard  and  yellow,  part  crystalline  and  part  (about  30 
per  cent)  amorphous.  Amorphous  sulphur  is  sometimes 
found  in  flowers  of  sulphur  and  can  be  detected  by  its  in- 
solubility in  carbon  disulphide.  Other  varieties  of  sulphur 
also  contain  amorphous  sulphur.  One  is  a  white  or  whitish 
powder,  made  by  boiling  flowers  of  sulphur  with  milk  of 
lime  and  adding  hydrochloric  acid  to  the  decanted  liquid ; 
a  fine  sulphur  powder  is  precipitated,  which  gives  the  liquid 
the  appearance  of  milk,  hence  the  name  often  applied  to  it, 
"  milk  of  sulphur." 

Uses.  —  Sulphur  is  used  in  making  sulphuric  acid  and 
other  sulphur  compounds,  gunpowder,  fireworks,  matches, 
in  vulcanizing  rubber,  and  as  a  medicine.  Considerable  is 
used  to  kill  phylloxera  (an  insect  which  destroys  grapevines) ; 
some  insecticides,  made  from  sulphur,  e.g.  lime-sulphur 
sprays,  contain  unstable  compounds  which  by  decomposi- 
tion liberate  sulphur  upon  the  insect  pest. 

The  Important  Compounds  of  Sulphur  are  hydrogen  sul- 
phide and  other  sulphides  (especially  metallic  sulphides), 
sulphur  dioxide  and  trioxide,  the  sulphites,  sulphuric  acid  and 
the  sulphates,  and  carbon  disulphide. 

HYDROGEN  SULPHIDE 

Hydrogen  sulphide,  H2S,  is  a  gaseous  compound  of  sulphur 
and  hydrogen,  and  is  sometimes  called  sulphuretted  hydrogen 
or  hydrosulphuric  acid.  It  occurs  in  some  volcanic  gases  and 
in  the  waters  of  sulphur  springs.  It  is  often  found  in  the  air, 
especially  near  sewers  and  cesspools,  since  it  is  one  product 
of  the  decay  of  organic  substances  containing  sulphur.  It 
is  one  of  the  impurities  of  crude  illuminating  gas. 


338  INORGANIC  CHEMISTRY 

Preparation.  —  The  gas  is  prepared  in  the  laboratory  by 
the  interaction  of  dilute  acids  and  metallic  sulphides;  usually 
dilute  hydrochloric  acid  and  ferrous  sulphide  are  used. 
When  the  acid  is  poured  upon  fragments  of  the  sulphide, 
the  gas  is  rapidly  evolved  without  applying  heat,  and  may 
be  collected  over  water.  The  equation  for  the  chemical 
change  is  — 

FeS     +      2HC1     =     H2S    +  FeCl2 

Iron  Hydrochloric        Hydrogen  Iron 

Sulphide  Acid  Sulphide         Chloride 

Properties.  —  Hydrogen  sulphide  gas  is  colorless  and  has 
the  odor  of  rotten  eggs.  It  is  poisonous.  A  little,  even  if 
diluted  with  air,  often  produces  headache  and  nausea,  and 
a  large  quantity  of  the  gas  is  fatal.  Care  should  be  used 
in  working  with  hydrogen  sulphide,  especially  if  the  genera- 
tor is  large.  The  dry  gas  is  slightly  heavier  than  air;  a 
liter  under  standard  conditions  weighs  1.537  gm.  It  has 
been  liquefied  and  solidified  by  the  usual  methods.  Hydro- 
gen sulphide  is  soluble  in  water,  one  volume  of  water  dis- 
solving about  three  volumes  of  the  gas  at  the  ordinary 
temperature;,  the  dissolved  gas  can  be  completely  removed 
by  boiling  the  solution.  The  solution  is  called  hydrogen 
sulphide  water,  and  is  often  used  instead  of  the  gas  in 
chemical  experiments.  The  solution  decomposes  slowly, 
sulphur  being  deposited. 

Hydrogen  sulphide  gas  is  inflammable  and  burns  with  a 
bluish  flame,  forming  water  and  sulphur  dioxide,  thus:  — 

2  H2S  +  3  O2  =  2  S02  +  2  H20 

Hydrogen       Oxygen        Sulphur         Water 
Sulphide  Dioxide 

If  the  supply  of  air  is  insufficient,  combustion  is  incom- 
plete, water  and  sulphur  being  formed,  thus :  — 

2H2S     +     O2     =     2H20     +     2S 

Hydrogen         Oxygen  Water  Sulphur 

Sulphide 


SULPHUR  AND   ITS   COMPOUNDS  339 

It  is  a  reducing  agent,  and  is  used  as  such  in  chemical  ex- 
periments. Even  sulphuric  acid  is  reduced  by  it,  thus :  — 

H2SO4     +     H2S     =    SO2     -f      S       +     2H2O 

Sulphuric          Hydrogen       Sulphur        Sulphur  Water 

Acid  Sulphide         Dioxide 

Hydrogen  sulphide  gas  dissolved  in  water  gives  a  solution 
which  has  a  feeble  acid  reaction,  is  neutralized  by  bases, 
and  forms  salts  called  sulphides.  It  dissociates  only  slightly, 
the  ions  being  chiefly  H+  and  HS~  ;  some  S~  "  ions  are  formed. 

Composition  of  Hydrogen  Sulphide  Gas.  —  When  metals  are  heated 
in  dry  hydrogen  sulphide,  metallic  sulphides  are  formed;  and  the 
volume  of  hydrogen  liberated  by  their  decomposition  is  the  same  as 
the  original  volume  of  gas  used.  Since  the  hydrogen  molecule  con- 
tains two  atoms  (H2),  there  must  be  two  atoms  of  hydrogen  in  the 
hydrogen  sulphide  molecule.  The  vapor  density  of  hydrogen  sul- 
phide gas  as  found  by  experiment  requires  the  molecular  weight 
34  (approximately).  Subtracting  2  for  2  H,  the  remainder  (32)  agrees 
well  with  the  atomic  weight  of  sulphur.  Hence,  there  can  be  only 
one  atom  of  sulphur  in  hydrogen  sulphide,  and  the  formula  must 
be  H2S. 

Sulphides  may  be  regarded  as  salts  of  the  weak  acid, 
hydrosulphuric  acid,  though  they  are  not  always  prepared 
directly  from  hydrogen  sulphide  by  the  substitution  of  a 
metal  for  its  hydrogen.  They  may  be  produced  by  the 
direct  union  of  sulphur  and  metals  (as  in  the  case  of  iron 
and  copper  sulphides  previously  mentioned),  by  exposing 
the  metal  to  the  moist  gas,  or  by  the  reduction  of  a  sul- 
phate with  carbon.  A  more  common  way  is  to  precipitate 
them  by  passing  the  gas  into  solutions  of  metallic  compounds, 
or  sometimes,  by  adding  hydrogen  sulphide  water  to  such 
solutions.  Copper,  tin,  lead,  and  silver  are  rapidly  tar- 
nished by  the  gas.  Silverware,  on  this  account,  turns  brown 
or  black,  especially  in  houses  heated  by  coal  or  lighted  by 
coal  gas.  The  brown  silver  sulphide  also  coats  silver  spoons 


340  INORGANIC   CHEMISTRY 

which  are  put  into  mustard  or  eggs.  Lead  compounds  are 
blackened  by  this  gas,  owing  to  the  formation-  of  lead  sul- 
phide, thus : — 

PbO    +     H2S     =     PbS     +     H2O 

Lead  Hydrogen  Lead  Water 

Oxide          Sulphide          Sulphide 

For  this  reason  houses  painted  with  "white  lead"  paint 
often  become  dark,  and,  similarly,  oil  paintings  are  dis- 
colored. The  darkening  (to  brown  or  black)  of  the  solu- 
tion of  a  lead  compound  is  the  customary  test  for  hydrogen 
sulphide. 

Many  sulphides  have  a  brilliant  color.  Arsenious  sulphide  is 
pale  yellow,  cadmium  sulphide  is  golden  yellow,  manganese  sulphide 
is  flesh  colored,  zinc  sulphide  is  white,  antimony  sulphide  is  orange 
red.  They  vary  in  solubility.  Most  sulphides  are  insoluble  in  water. 
The  sulphides  of  lead,  silver,  copper,  and  some  other  metals  are  in- 
soluble in  dilute  hydrochloric  acid.  The  sulphides  of  iron,  zinc,  and 
some  other  metals  are  decomposed  by  dilute  hydrochloric  acid,  but 
are  precipitated  if  ammonium  hydroxide  is  present.  Sulphides  of 
certain  metals  dissolve  in  water.  Hence  by  precipitating  metals 
under  different  conditions,  groups  of  metals  may  be  separated  and 
subjected  to  further  tests.  The  color  often  affords  a  ready  means  of 
detecting  each  sulphide.  Hydrogen  sulphide  is  thus  a  serviceable 
reagent  in  Qualitative  Analysis. 

OXIDES   OF  SULPHUR 

Sulphur  Dioxide,  SO2,  is  the  common  compound  of  sulphur 
and  oxygen.  It  occurs  in  the  gases  of  volcanoes,  and  to  a 
slight  extent  in  the  atmosphere,  since  it  is  the  usual  product 
of  the  combustion  of  sulphur  and  sulphur  compounds. 

Preparation.  —  When  sulphur  burns  in  air  (or  oxygen), 
sulphur  dioxide  is  formed,  thus :  — 

S       +      O2     =      SO2 

Sulphur         Oxygen          Sulphur 
Dioxide 


SULPHUR   AND   ITS   COMPOUNDS  341 

It  is  also  formed  by  roasting  iron  disulphide  (iron  pyrites) 
in  the  air,  thus :  — 

4  FeS2     +     11  O2  =  8  S02  +  2  Fe2O3 

Iron  Oxygen        Sulphur  Iron 

Disulphide  Dioxide  Oxide 

The  foregoing  reaction  is  utilized  on  a  large  scale  in  the 
commercial  manufacture  of  sulphuric  acid.  Sulphur  and 
carbon  reduce  sulphuric  acid  to  sulphur  dioxide,  thus:  — 

S     +     2H2SO4  =  3S02   +    2H2O 

Sulphur         Sulphuric          Sulphur 
Acid  Dioxide 

C    +    2  H2SO4  =  2  S02  .+  CO2  +  2  H2O 

Carbon  Carbon 

Dioxide 

Two  methods  of  preparation  are  used  in  the  laboratory. 

(1)  If  copper  and  concentrated  sulphuric  acid  are  heated, 
a  series  of  complex  changes  results  finally  in  the  evolution 
pf  sulphur  dioxide.     The  equation  is  usually  written :  — 

Cu  +  2H2SO4  =  SO2  +   CuSO4   +    2  H2O 

Copper        Sulphuric        Sulphur        Copper 
Acid  Dioxide       Sulphate 

(2)  Sulphuric  (or  hydrochloric)  acid  dropped  upon   a   sul- 
phite yields  sulphur  dioxide,  thus :  — 

Na2SO3  +  H2SO4  =    SO2    +  Na2SO4  +  H2O 

Sodium         Sulphuric       Sulphur          Sodium 
Sulphite  Acid  Dioxide         Sulphate 

The  latter  method  is  convenient  for  liberating  a  steady 
current  of  the  gas.  (See  page  344.) 

Properties.  —  Sulphur  dioxide  gas  has  no  color.  It  has 
a  suffocating  odor,  being  the  well-known  odor  readily  noticed 
when  sulphur  is  burned.  It  will  not  burn  in  the  air,  nor 
will  it  support  ordinary  combustion.  A  burning  taper  or 
stick  of  wood  is  instantly  extinguished  by  it,  but  finely 


342  INORGANIC  CHEMISTRY 

divided  metals,  iron  for  example,  burn  in  it.  It  is  a  heavy 
gas,  the  high  density  (2.2)  allowing  it  to  be  readily  collected 
by  displacement  of  air,  as  in  the  case  of  chlorine.  A  liter 
of  sulphur  dioxide  under  standard  conditions  weighs  2.927 
gm.  Its  critical  temperature  is  about  +  155.5°  C.,  and  at 
a  moderately  low  temperature  it  changes  into  a  transparent, 
colorless  liquid,  which  boils  at  —  8°  C.;  if  sufficiently  cooled, 
it  freezes  into  a  transparent,  icelike  solid,  which  melts  at 
about  —  76°  C.  Liquid  sulphur  dioxide  is  a  common  com- 
mercial article.  Sulphur  dioxide  gas  is  very  soluble  in 
water.  At  the  ordinary  temperature  one  volume  of  water 
dissolves  about  forty  volumes  of  gas,  but  the  solution  loses 
it  all  by  boiling.  This  solution  is  sour  and  reddens  blue 
litmus,  and  contains  sulphurous  acid  (see  below).  Moist 
sulphur  dioxide  bleaches  vegetable  coloring  matters.  A  red 
or  purple  flower  loses  color  in  it.  Silk,  hair,  straw,  wool, 
and  other  delicate  substances,  which  would  be  injured  by 
chlorine,  are  whitened  by  sulphur  dioxide.  In  some  cases 
the  color  returns  when  the  bleached  article  is  exposed  to 
the  air  for  some  time,  and  usually  such  bleached  objects 
become  yellow  with  age.  The  coloring  matter  is  not  wholly 
destroyed,  but  probably  unites  with  the  sulphur  dioxide  to 
form  a  colorless  compound,  which  slowly  decomposes. 

Uses.  —  Immense  quantities  of  sulphur  dioxide  are  used 
in  the  manufacture  of  sulphuric  acid.  The  gas  is  also  used 
to  preserve  meat  and  wines,  to  fumigate  clothing  and  houses, 
in  paper  making,  in  tanning,  in  refining  sugar,  and  in  mak- 
ing acid  sodium  sulphite.  Liquid  sulphur  dioxide  is  used 
in  extracting  glue  and  gelatine,  and  in  various  metallurgical 
processes.  It  absorbs  considerable  heat  during  evapora- 
tion, and  was  formerly  used  in  some  ice  machines. 

The  Composition  of  Sulphur  Dioxide  is  based  on  the  following: 
The  gas  formed  by  burning  sulphur  in  a  measured  volume  of  oxygen 


SULPHUR  AND   ITS   COMPOUNDS  343 

has  the  same  volume  as  the  oxygen  itself.  Hence  there  are  as  many 
molecules  of  sulphur  dioxide  as  there  were  of  oxygen;  that  is,  one 
molecule  of  sulphur  dioxide  contains  two  atoms  of  oxygen.  Two 
atoms  of  oxygen  weigh  32.  But  the  molecular  weight  of  sulphur 
dioxide  calculated  from  an  experimental  determination  of  its  vapor 
density  is  about  64.  Subtracting  32  (i.e.  2x  16)  for  the  oxygen  (2  O) 
from  this,  there  remains  about  32  for  sulphur.  The  atomic  weight 
of  sulphur  is  32.07;  hence  sulphur  dioxide  contains  only  one  atom 
of  sulphur,  and  its  composition  is  expressed  by  the  formula  SO2. 

Sulphur  Trioxide,  SO3,  is  formed  by  the  direct  union  of 
sulphur  dioxide  and  oxygen,  though  a  little  is  produced 
when  sulphur  burns  in  air  or  in  oxygen.  The  action  is  slow, 
but  can  be  hastened  by  passing  a  mixture  of  sulphur  dioxide 
and  oxygen  (or  air)  over  hot  platinum,  or  better  over  asbestos 
coated  with  platinum.  Other  substances  also  hasten  the 
chemical  change.  At  the  ordinary  temperature  sulphur 
trioxide  is  a  liquid,  which  boils  at  46°  C.  and  solidifies  at 
15°  C.  to  a  white  crystalline  mass.  Another  solid  form 
(sulphur  hexoxide,  S206),  resembling  asbestos,  is  known. 
When  exposed  to  moist  air,  it  fumes  strongly,  forming  sul- 
phuric acid  ;  and  when  dropped  into  water,  it  dissolves  with 
a  hissing  sound  and  evolution  of  heat.  The  equation  for  this 
chemical  change  is  — 

SO3    +    H,O  =  H2SO4 

Sulphur         Water        Sulphuric 
Trioxide  Acid 

OXYGEN  ACIDS  AND  SALTS  OF  SULPHUR 

Sulphurous  Acid  and  Sulphites.  —  Sulphurous  acid,  H2SO3, 
is  formed  when  sulphur  dioxide  dissolves  in  water.  Sulphur 
dioxide  is,  therefore,  sulphurous  anhydride.  The  simplest 
equation  expressing  this  fact  is  — 

S02   +   H20  =   H2S03 

Sulphur       Water        Sulphurous 
Dioxide  Acid 


344  INORGANIC  CHEMISTRY 

This  acid  is  known  only  in  solution,  and  resembles  car- 
bonic acid  in  this  respect.  It  is  unstable  and  decomposes 
readily  into  sulphur  dioxide  and  water.  It  gradually  forms 
sulphuric  acid  by  combining  with  oxygen  from  the  air  and 
very  rapidly  by  interaction  with  oxidizing  agents,  such 
as  potassium  permanganate.  Sulphurous  acid  is  dibasic, 
and  forms  two  classes  of  salts,  the  normal  and  acid  sul- 
phites. They  are  reducing  agents,  and  yield  sulphur  dioxide 
when  treated  with  acids.  It  will  be  recalled  in  this  con- 
nection that  a  convenient  method  of  preparing  sulphur 
dioxide  is  the  decomposition  of  normal  sodium  sulphite 
(NajjSOs)  by  sulphuric  or  hydrochloric  acid.  The  sulphite 
first  forms  sulphurous  acid  (H2S03),  and  the  unstable  acid 
decomposes  into  water  and  sulphur  dioxide.  The  following 
equations  express  these  two  reactions  :  — 

Na2S03  +  H2SO4   =    H2S03   +   Na2SO4 

Sodium        Sulphuric        Sulphurous          Sodium 
Sulphite  Acid  Acid  Sulphate 

H2SO3     =     S02    +    H20 

Sulphurous          Sulphur        Water 
Acid  Dioxide 

Acid  sodium  sulphite  (HNaSO3),  often  called  bisulphite  of 
soda,  is  the  anticolor  used  to  remove  the  excess  of  chlorine 
from  bleached  cotton  cloth.  It  is  also  used  in  brewing, 
tanning,  and  in  the  manufacture  of  starch,  sugar,  and 
paper.  Acid  calcium  sulphite  (CaH2(SO3)2),  also  called 
bisulphite  of  calcium,  is  prepared  by  passing  sulphur  di- 
oxide into  milk  of  lime,  and  is  used  in  paper  making. 

Sulphurous  acid  is  rather  weak.  In  terms  of  the  theory  of 
electrolytic  dissociation  it  dissociates  only  to  a  compara- 
tively slight  extent,  the  ions  being  chiefly  H+  and  HSO3~; 
the  latter  ion,  however,  dissociates  to  some  degree  into  H+ 
and  SO3~~.  Solutions  of  the  acid  sulphites  of  sodium  and 
potassium  have  an  acid  reaction  owing  to  the  slight  disso- 


SULPHUR  AND   ITS   COMPOUNDS  345 

elation  of  the  HS03-ion  into  the  ions  H+  and  SO3~~.  Solu- 
tions of  the  normal  sulphites  of  sodium  and  potassium  are 
alkaline.  This  fact  is  briefly  explained  in  terms  of  the 
theory  of  electrolytic  dissociation  as  follows :  The  SO3-ions 
from  the  sulphite  tend  to  form  the  more  stable  HS03-ions 
by  combining  with  the  H-ions  furnished  by  the  appreciably 
dissociated  water;  hence  OH-ions  are  left  in  the  solution 
and  give  it  an  alkaline  reaction.  As  already  stated,  cases 
of  double  decomposition  like  the  one  just  cited  in  which 
water  is  one  factor  are  illustrations  of  hydrolysis  (see 
Hydrolysis,  Chapter  X). 

Sulphuric  Acid,  H2SO4,  is  found  in  the  waters  of  a  few 
rivers  and  mineral  springs.  It  is  manufactured  in  enor- 
mous quantities  and  used  for  many  purposes. 

Sulphuric  acid  was  doubtless  known  to  the  Arabian  alchemists 
living  in  the  tenth  century.  It  was  definitely  mentioned  by  Basil 
Valentine  in  the  fifteenth  century,  who  described  its  preparation 
by  heating  a  mixture  of  iron  sulphate  (green  vitriol)  and  sand.  The 
product,  an  oily  liquid,  was  called  oil  of  vitriol,  a  name  now  often 
applied  to  commercial  sulphuric  acid. 

Sulphuric  acid  is  manufactured  by  two  processes,  the  lead 
chamber  process  and  the  contact  process.  In  both  processes 
sulphur  dioxide  is  oxidized  to  sulphur  trioxide,  which  is 
transformed  into  sulphuric  acid  by  combination  with  water. 
A  skeleton  equation,  so  to  speak,  may  be  written  thus : 

2SO2       +       O2       +       2H2O  2H2SO4 

Sulphur  Oxygen  Water  Sulphuric 

Dioxide  Acid 

The  oxidation  is  accomplished  in  two  ways.  In  the  older 
or  lead  chamber  process  oxides  of  nitrogen  are  used.  In 
the  newer  or  contact  process  a  suitable  catalytic  agent,  usually 
finely  divided  platinum,  is  employed. 


346  INORGANIC  CHEMISTRY 

In  the  chamber  process  sulphur  dioxide,  air,  steam,  and 
oxides  of  nitrogen  are  introduced  into  large  lead  chambers. 
These  gases  interact  and  produce  sulphuric  acid,  which  col- 
lects on  the  floors  of  the  lead  chambers.  The  chemical 
changes  involved  in  this  process  are  complex  and  variable. 
The  main  reactions  are  :  — 

2  SO2  +  N2O3  +  O2  +  H2O  -  =  2  SO2(OH)(ONO) 

Nitrosyl-sulphuric 
Acid 

2  S02(OH)(ONO)  +  H2O  =  2  H2SO4  +  N203 

The  nitrogen  trioxide  (N2O3)  is  apparently  an  essential 
factor,  though  according  to  some  authorities  the  change 
may  be  due  to  a  mixture  of  nitric  oxide  (NO)  and  nitrogen 
peroxide  (NO2).  Theoretically  a  small  quantity  of  oxides 
of  nitrogen  is  needed  to  form  an  unlimited  amount  of  sul- 
phuric acid.  Some  is  lost,  however,  and  must  be  replaced. 
This  is  done  by  putting  nitric  acid  into  the  top  of  the  Glover 
tower  or  by  injecting  it  into  the  chambers. 

The  details  of  the  lead  chamber  apparatus  are  shown  in  Figure 
53.  There  are  three  main  parts :  (a)  the  furnace  for  producing 
sulphur  dioxide,  (6)  the  lead  chambers  together  with  the  Glover 
and  Gay-Lussac  towers  for  changing  the  gaseous  mixture  into  sul- 
phuric acid,  and  (c)  the  concentrating  apparatus.  The  manufac- 
ture is  conducted  as  follows:  (1)  Sulphur  or  iron  disulphide  (iron 
pyrites,  FeS2)  is  burned  in  a  furnace  with  enough  air  to  change  the 
sulphur  into  sulphur  dioxide,  and  to  furnish  the  proper  amount  of 
oxygen  for  later  changes.  In  some  works  the  furnace  is  provided 
with  "  niter  pots  "  to  produce  nitric  acid  vapor.  (2)  The  mixture 
of  sulphur  dioxide,  oxides  of  nitrogen,  and  air  passes  from  the  fur- 
nace into  the  bottom  of  the  Glover  tower.  This  is  a  tall  tower 
filled  with  small  stones  over  which  flow  two  streams  of  sulphuric 
acid,  one  dilute  and  the  other  containing  nitrogen  dioxide  (from 
the  Gay-Lussac  tower  acid).  These  acids  cool  the  ascending  gases  ; 
the  dilute  acid  is  deprived  of  water  and  the  tower  acid  of  nitro- 
gen dioxide.  Hence,  concentrated  acid  flows  out  of  the  bottom  of 


SULPHUR  AND  ITS  COMPOUNDS 


the  Glover  tower, 
while  from  the 
top  sulphur  diox- 
ide, oxides  of  ni- 
trogen, steam,  and 
air  pass  into  the 
first  lead  chamber. 
Here  nitric  acid 
is  introduced,  as 
well  as  steam. 
The  main  chemi- 
cal changes  occur 
in  this  and  in  the 
second  chamber. 
These  chambers 
are  huge  boxes 
of  sheet  lead 
supported  on  a 
wooden  frame- 
work. The  un- 
used gases  pass  on 
into  the  bottom 
of  the  Gay-Lussac 
tower.  This  tower 
is  filled  with  coke 
or  earthenware 
over  which  flows 
concentrated  sul- 
phuric acid  (ob- 
tained from  the 
Glover  tower), 
which  absorbs  the 
unused  nitrogen 
dioxide.  The  ox- 
ide,  as  stated 
above,  is  liber- 
ated again  in  the 
Glover  tower. 
(3)  The  acid  pro- 
duced in  the  cham- 


FIG.  53. — Apparatus  for  chamber  process. 


INORGANIC  CHEMISTRY 


bers  contains  from  60  to  70  per  cent 
of  H2SO4.  It  is  concentrated  into 
commercial  sulphuric  acid,  which 
contains  about  96  to  98  per  cent,  by 
evaporation,  first  in  lead-lined  pans 
and  finally  in  a  platinum  vessel. 

In  the  contact  process  sul- 
phur dioxide  and  air,  carefully 
purified  and  heated  to  about 
400°  C.,  are  brought  in  contact 
with  a  catalytic  agent  (see  page 
185),  which  is  usually  finely 
divided  platinum.  The  sulphur 
dioxide  is  oxidized  to  sulphur 
trioxide,  thus :  — 

2  SO2  +  O2  =  2  SO3 

The  sulphur  trioxide  is  con- 
ducted into  sulphuric  acid  con- 
taining a  little  water  with  which 
the  trioxide  combines,  thereby 
forming  sulphuric  acid,  thus  :  — 

SO3  +  H2O  =  H2SO4 

The  contact  process  apparatus 
is  shown  in  Figure  54.  The  blower 
A  forces  air  into  the  burner  B,  where 
the  suphur  dioxide  is  formed  by 
burning  iron  pyrites  (FeS2)  or  sul- 
phur. The  gases  pass  into  C,  where 
they  are  freed  from  sulphur  dust 
and  other  solid  impurities.  The 
gases,  cooletl  by  the  pipe  D,  are 
further  cleaned  in  the  scrubbers, 
which  contain  coke  wet  with  water 
(E)  and  with  sulphuric  acid  (F), 


SULPHUR  AND  ITS  COMPOUNDS  349 

and  then  freed  from  arsenic  compounds  in  G.  Traces  of  such  com- 
pounds "  poison  "  the  platinum  and  stop  the  formation  of  sulphur 
trioxide.  The  purified  gases  (mainly  sulphur  dioxide)  then  enter 
the  mixer  and  heater  H.  Here  a  large  excess  of  air  is  introduced 
and  the  whole  mixture  is  heated  to  about  400°  C.  The  mixture  of 
sulphur  dioxide  and  air  next  passes  into  the  contact  chamber  /. 
Here  the  gases  come  in  contact  with  the  catalytic  agent  and  form 
sulphur  trioxide.  The  catalytic  agent,  if  platinum,  consists  of 
asbestos  fibers  coated  with  a  very  thin  layer  of  metallic  platinum 
and  is  spread  out  on  plates  or  mixed  with  porous  material  in  order 
to  provide  a  large  contact  surface.  The-  sulphur  trioxide  finally 
passes  into  an  absorber  J  partly  filled  with  sulphuric  acid  contain- 
ing 2  to  3  per  cent  of  water.  In  this  liquid,  the  combination  with 
water  takes  place  readily,  water  being  added  to  maintain  the  required 
concentration  in  the  absorber. 

Properties  of  Sulphuric  Acid.  —  Sulphuric  acid  is  an  oily 
liquid,  colorless  when  pure,  though  often  brown  from  the 
presence  of  charred  organic  matter,  such  as  dust  and  straw. 
The  specific  gravity  of  the  commercial  acid  is  about  1.83. 
When  sulphuric  acid  is  mixed  with  water,  much  heat  is 
evolved.  The  acid  should  always  be  poured  into  the  water, 
otherwise  the  intense  heat  may  crack  the  vessel  or  spatter 
the  hot  acid.  Sulphuric  acid  combines  with  water  and  forms 
hydrates ;  a  rather  stable  one  has  the  formula :  — 

H2SO4 .  H2O. 

The  tendency  to  absorb  water  is  shown  in  many  ways. 
The  concentrated  acid  absorbs  moisture  from  the  air  and 
from  gases  passed  through  it.  It  is  often  used  in  the  labora- 
tory to  dry  gases,  since  it  is  not  volatile  at  the  ordinary 
temperature.  Wood,  paper,  sugar,  starch,  cotton  cloth, 
and  many  organic  substances  are  blackened  by  sulphuric 
acid.  Such  compounds  contain  hydrogen  and  oxygen  in 
the  proportion  to  form  water;  these  two  elements  are  ab- 
stracted and  carbon  alone  remains.  Similarly,  sulphuric 


350  INORGANIC  CHEMISTRY 

acid  withdraws  water  from  the  flesh,  making  painful 
wounds. 

Sulphuric  acid  boils  at  about  338°  C.,  but  begins  to  give 
off  fumes  of  sulphur  trioxide  at  about  150°  C.  It  oxidizes 
carbon  and  sulphur  to  carbon  dioxide  and  sulphur  dioxide, 
owing  to  its  instability  when  heated  with  these  elements. 
Metals  decompose  it,  yielding  various  products,  such  as 
hydrogen,  sulphur  dioxide,  or  hydrogen  sulphide.  It  com- 
bines directly  with  ammonia  gas,  thus  :  — 

2NH3      +       H2S04       =       (NH4)2S04 

Ammonia  Sulphuric  Ammonium 

Gas  Acid  Sulphate 

Its  interaction  with  salts,  such  as  chlorides,  nitrates,  and 
sulphites,  which  results  in  the  liberation  of  the  correspond- 
ing acid,  has  already  been  discussed.  A  solution  of  sul- 
phuric acid  in  water  contains  hydrogen  ions  (H+),  monova- 
lent  ions  (HSO4~),  and  divalent  ions  (SO4~~),  depending  upon 
the  concentration.  Dilute  solutions  contain  an  abundance 
of  SO4-ions  (S04— )  and  H-ions  (H+). 

Uses  of  Sulphuric  Acid.  —  Sulphuric  acid  is  one  of  the 
most  important  substances.  Directly  or  indirectly  it  is 
used  in  hundreds  of  industries  upon  which  the  comfort, 
prosperity,  and  progress  of  mankind  depend.  It  is  used 
in  the  manufacture  of  other  mineral  acids  and  many  organic 
acids.  It  is  essential  in  one  process  of  manufacturing 
sodium  carbonate.  Enormous  quantities  are  consumed  in 
making  artificial  fertilizers,  alum,  nitroglycerin,  glucose, 
phosphorus,  dyestuffs,  and  in  various  parts  of  such  funda- 
mental industries  as  dyeing,  bleaching,  electroplating,  refin- 
ing, and  metallurgy. 

Sulphates.  —  Sulphuric  acid  is  dibasic  and  forms  two 
classes  of  salts,  —  the  normal  sulphates,  such  as 


SULPHUR  AND   ITS   COMPOUNDS  351 

and  the  acid  sulphates,  such  as  HNaS04.  The  normal  sul« 
phates  are  rather  stable  salts ;  many  yield  sulphur  trioxide 
when  heated  to  a  high  temperature.  The  acid  salts  lose 
water  when  heated,  yielding  in  addition  salts  called  pyro- 
sulphates  (see  below  under  pyrosulphuric  acid).  Most  sul- 
phates are  soluble  in  water,  only  the  sulphates  of  barium, 
strontium,  and  lead  being  insoluble,  while  calcium  sulphate 
is '  slightly  soluble.  Important  sulphates  are  calcium  sul- 
phate (gypsum,  CaSO4.2H2O),  barium  sulphate  (heavy 
spar,  BaSO4),  zinc  sulphate  (white  vitriol,  ZnSOJ,  copper 
sulphate  (blue  vitriol  or  bluestone,  CuSO*),  iron  sulphate 
(green  vitriol,  copperas,  ferrous  sulphate,  FeSOJ,  sodium 
sulphate  (Glauber's  salt,  Na2SO4),  and  magnesium  sulphate 
(Epsom  salts,  MgSO4,  and  kieserite,  MgSO4.H2O).  Sul- 
phates are  widely  used  as  medicine  and  in  many  industries. 
They  are  described  more  fully  under  the  individual  metals. 

The  test  for  sulphuric  acid  or  a  soluble  sulphate  is  the 
formation  of  white,  insoluble  barium  sulphate  upon  the 
addition  of  barium  chloride  solution.  An'insoluble  sulphate, 
such  as  calcium  sulphate,  when  fused  on  charcoal  is  reduced 
to  a  sulphide,  which  blackens  a  moist  silver  coin.  The 
blackening  is  due  to  the  formation  of  silver  sulphide. 

The  usual  test,  as  already  stated,  is  applicable  to  both  the 
acid  and  soluble  sulphates  because  their  solutions  contain 
SO4-ions.  When  a  solution  containing  barium  ions  is  added, 
the  ions  Ba++  and  SO4  unite  and  form  barium  sulphate, 
which  is  insoluble  in  water  and  is  not  (like  some  other 
barium  salts)  decomposed  by  the  common  acids.  The  ionic 
equation  for  the  reaction  is  :  — 

Ba+  +  +  S04-  -  — >-  BaSO4 

Fuming  Sulphuric  Acid,  H2S2O7,  is  made  by  adding  sul- 
phur trioxide  to  sulphuric  acid,  or  by  heating  moist  ferrous 
sulphate.  It  is  sometimes  called  Nordhausen  sulphuric 


352  INORGANIC  CHEMISTRY 

acid.  It  is  a  thick,  brown  liquid,  whiclr  fumes  strongly  in 
the  air,  owing  to  the  escape  of  oxides  of  sulphur.  It  is 
used  in  gas  analysis  to  absorb  ethylene  and  other  illuminants, 
and  in  dyeing  to  dissolve  indigo.  If  the  fuming  acid  is 
cooled  to  0°  C.,  crystals  separate;  they  are  called  pyrosul- 
phuric  acid  or  disulphuric  acid.  Its  salts  are  the  pyro- 
sulphates. 

OTHER  SULPHUR  COMPOUNDS 

Sodium  Thiosulphate,  Na2S2O3;  is  a  salt  of  an  unstable 
acid.  It  is  sometimes  incorrectly  called  sodium  hyposul- 
phite, or  simply  "hypo."  It  is  a  white,  crystalline  solid, 
very  soluble  in  water.  The  solution,  used  in  excess,  dis- 
solves the  halogen  compounds  of  silver,  i.e.  AgCl,  AgBr, 
Agl;  hence  its  extensive  use  in  photography.  (See  Photog- 
raphy.) It  is  used  in  dyeing,  and  in  chemical  analysis 
(for  determining  the  amount  of  free  iodine  in  a  solution). 

Carbon  Bisulphide  (or  Carbon  Bisulphide,  CS2)  when  pure 
is  a  clear,  colorless,  liquid,  with  an  agreeable  odor;  the  com- 
mercial substance  is  yellow  and  has  an  offensive  odor.  It  is 
poisonous.  Carbon  disulphide  is  an  endothermic  com- 
pound, i.e.  its  synthesis  from  its  elements  is  accompanied  by 
absorption  of  heat.  The  thermal  equation  for  the  forma- 
tion of  carbon  disulphide  is  — 

C      +     S2        =          CS2  28,700  cal. 

Carbon          Sulphur  Carbon 

Disulphide 

Like  all  endothermic  compounds  it  is  relatively  unstable 
and  can  be  exploded  by  mercuric  fulminate.  Ordinarily  it 
can  be  handled  without  danger  of  explosion.  It  is  volatile 
and  extremely  inflammable,  the  equation  for  its  combustion 
being  — 

CS2       +       3  02       =       CO2       +       2  SO2 

Carbon  Oxygen  Carbon  Sulphur 

Disulphide  Dioxide  Dioxide 


SULPHUR   AND    ITS   COMPOUNDS 


353 


This  liquid  is  practically  insoluble  in  water.  It  dissolves 
rubber,  gums,  fats,  resins,  iodine,  camphor,  and  some  forms 
of  sulphur.  It  is  a  highly  refracting  liquid,  and  hollow 
glass  prisms  filled  with  it  are  used  to  decompose  light.  As 
a  solvent  it  is  used  to  dissolve  pure  rubber  in  the  manu- 
facture of  rubber  cement.  It  is  also  used  to  kill  insects  on 
both  living  and  dried  plants  (e.g.  in  museums),  and  to  ex- 
terminate burrowing  animals,  such  as  moles  and  wood- 
chucks.  Many  oils,  waxes,  and 
greases  are  extracted  by  carbon 
disulphide. 

Until  recently  carbon  disulphide 
was  manufactured  by  passing  sul- 
phur vapor  over  red-hot  coke  or 
charcoal  in  iron  or  earthenware 
retorts.  The  product  required 
laborious  purification.  It  is  now 
manufactured  by  an  electrothermal 
process  in  a  furnace  somewhat  like 
that  shown  in  Fig.  55.  Several 
groups  of  carbon  electrodes  (E,  E) 
are  set  into  the  base  of  a  furnace, 
coke  is  packed  loosely  between  them; 

Sulphur    is    put    into   Z     and   partly    FIG.    55. —  Furnace    for    the 

surrounds  the  electrodes.  The  body  manufacture  of  carbon  di- 
of  the  furnace  is  filled  with  char-  sulPhide- 
coal  (C).  Sulphur  is  introduced  at  suitable  points  (S,  S,  S), 
and  coke  is  fed  in  through  K,  K.  When  the  current  passes, 
the  sulphur  melts  and  unites  with  the  heated  carbon  above 
the  electrodes.  The  vapors  of  the  resulting  carbon  disulphide 
escape  from  the  top  of  the  furnace  through  P,  and  are 
condensed  in  a  special  apparatus. 


354  INORGANIC  CHEMISTRY 

f 

PROBLEMS  AND  EXERCISES 

1.  Sulphur  is  burned  in  10  1.  of  air  containing  21  per  cent  of 
oxygen  by  volume  and  measured  at  22°  C.  and  767  mm.     Calculate 
(a)  the  weight  of  sulphur  burned,  (6)  the  volume  of  oxygen  con- 
sumed, (c)  the  weight  of  sulphur  dioxide  produced. 

2.  Suppose  a  manufacturer  of  H2SO4  starts  with  100  tons  of 
sulphur  and  obtains  the  theoretical  yield  in  each  case:    (a)  What 
weight  of  oxygen  is  needed  to  burn  the  sulphur  to  sulphur  dioxide  ? 
(6)  What  additional  weight  of  oxygen  to  convert  the  sulphur  dioxide 
to  sulphur  trioxide  ?     (c)  What  weight  of  water  to  convert  the  sul- 
phur trioxide  to  sulphuric  acid  ?     (d)  And  how  much  sulphuric  acid 
is  obtained  ? 

3.  What  weight  of  pure  H2SO4  can  be  manufactured  from  150 
tons  (2000  Ib.  each)  of  iron  pyrite  containing  92  per  cent  of  FeS2? 

4.  What  weight  of  sulphuric  acid  (sp.  gr.  1.8354)  is  contained 
in  a  cylindrical  tank  20  m.  long  and  1.3  m.  in  diameter? 

5.  Calculate  the  atomic  weight  of  sulphur  from  the  following 
data:  (a)  In  BaSO4,  Ba  =  58.85,  8  =  13.73,  O  =  27.42;  (6)  10  gm.  of 
silver  yield  11.4815  gm.  of  silver  sulphide;    (c)  2  gm.  of  lead  yield 
2.9284  gm.  of  lead  sulphate.     (Use  exact  atomic  weights.) 

6.  What  weight  of  carbon  disulphide  is  formed  by  the  inter- 
action of  (a)  carbon  and  200  kg.  of  sulphur  and  (6)  sulphur  and  200 
kg.  of  carbon? 

7.  Express   the  following  reactions  by  volumetric   equations: 
(a)  Sulphur  dioxide  and  oxygen  form  sulphur  trioxide,  (6)  hydrogen 
sulphide  and   oxygen  form   sulphur  dioxide  and   water    (vapor), 
(c)  carbon  disulphide  (vapor)  and  oxygen  form  sulphur  dioxide  and 
carbon  dioxide. 

8.  Write  the  formulas  of  (a)  the  acid  sulphates,  (6)  the  sulphites, 
(c)  the  acid  sulphites,  and  (d)  the  sulphates  of  Ba,  Ca,  K,  Na. 

9.  Write  the  formulas  of  the  sulphides  of  NH4,  Al,  Ba,  Ca,  Cr, 
Cu(ic),  Hg(ic),  Ag,  Sn(ous),  Zn. 

10.  Complete  and  balance  the  following  equations:    (a)  H2SO4 
+  H2S  =  S02  +  -      -  +  H2O ;       (6)     SrCO3  +  H2SO4  =  SrSO4  + 

+  H2O  ;      (c)    CdCl2  +  H2S  =  HC1  +  -    -;      (d)    Pb(NO3)2 

+  -      -  =  PbS04  +  HN03 ;    (e)  NH3  +  -         =  (NH4)2SO4. 

11.  If  a  sample  of  carbon  disulphide  has  a  specific  gravity  of 
1.25,  what  volume  can  be  made  from  2  metric  tons  of  pure  sulphur? 

12.  Calculate  the  solubility  product  of  lead  sulphate  if  the  ioni- 
zation  is  100  per  cent  and  the  molar  solubility  is  .00013. 


CHAPTER  XIX 
Classification  of  the  Elements 

Introduction.  —  In  the  preceding  chapters  emphasis  has 
been  laid  on  the  individual  elements,  especially  oxygen, 
hydrogen,  nitrogen,  carbon,  chlorine,  and  sulphur.  These 
elements  differ  from  each  other  in  many  ways,  and  if  this 
diversity  prevailed  among  all  the  eighty  elements,  it  would 
be  difficult  to  proceed  very  far  with  the  study  of  chemistry. 
Fortunately  the  elements  are  not  independent.  Certain 
ones  are  so  similar  in  their  chemical  relations  that  they  can 
be  put  into  the  same  class.  Several  characteristic  classes 
have  been  formed.  The  arrangement  into  groups  or  classes 
not  only  simplifies  the  study  of  the  elements,  but  reveals 
many  fundamental  relations. 

Classification  of  the  Elements.  —  Many  attempts  have  been 
made  to  classify  the  elements.  About  the  time  of  Lavoisier 
(1743-1794)  they  were  roughly  divided  into  metals  and  non- 
metals.  Those  elements  were  called  metals  which  were 
hard,  lustrous,  heavy,  and  good  conductors  of  heat,  while 
the  others  were  called  non-metals.  This  classification 
proved  to  be  misleading  as  additional  elements  were  dis- 
covered whose  properties  did  not  harmonize  with  the  prin- 
ciple of  division.  It  is  used  at  the  present  time,  however, 
because  many  common  elements  fall  readily  into  one  of 
these  classes,  as  shown  in  the  lists  of  metals  and  non-metals 
given  in  Chapter  X. 

Classification  according  to  acid  and  basic  properties  pre- 
vailed for  a  time.  .But  it  was  abandoned  largely  because 
such  a  basis  of  division  excluded  elements  exhibiting  both 

355 


356  INORGANIC   CHEMISTRY 

acid  and  basic  properties,  such  as  zinc  and  chromium.  The 
elements  have  also  been  classified  according  to  their  valence 
into  six  or  seven  groups  (the  mono-,  di-,  tri-,  etc.).  But 
this  plan  was  given  up  partly  because  of  so  many  cases  of 
variable  valence,  but  mainly  on  account  of  the  indiscrim- 
inate character  of  the  groups.  For  example,  elements  so 
unlike  as  sodium  and  chlorine  were  in  the  same  group. 

About  1828  striking  resemblances  between  certain  ele- 
ments were  pointed  out,  and  several  groups  or  families 
were  suggested.  For  example  :  — 

Lithium  Selenium  Calcium  Nitrogen 

Sodium  Sulphur  Strontium  Phosphorus 

Potassium  Oxygen  Barium  Arsenic 

This  classification  was  arbitrarily  based  on  selected  physical 
and  chemical  properties.  It  was  interesting  but  incom- 
plete, because  it  emphasized  resemblances  and  overlooked 
differences;  that  is,  the  basis  of  comparison  was  not  broad 
enough. 

The  first  actual  progress  began  to  be  made  about  1850, 
when  chemists  became  deeply  interested  in  the  significance 
of  atomic  weights.  Dumas  and  others  pointed  out  certain 
striking  numerical  relations  between  the  atomic  weights  of 
related  elements.  Thus,  the  atomic  weight  of  sodium  is 
half  the  sum  of  the  atomic  weights  of  lithium  and  potassium. 

Li  =  7,  Na  =  23,  K  =  39.     l±®>  =  23 
The  same  is  true  of  phosphorus,  arsenic,  and  antimony. 

P  =  31,  As  =  75,  Sb  =  120.     31  +  12Q  =  75.5 

2i 

A  number  of  such  families  was  found,  but  this  method  of 
classification  was  not  comprehensive,  nor  did  it  entirely 
eliminate  chance  selections. 

The   existence   of   other   relations   similar   to   those   just 


MENDELEJEFF 
1834-1907 


CLASSIFICATION    OF   THE   ELEMENTS  357 

cited,  a  deep  desire  to  obtain  more  accurate  atomic  weights, 
and  a  growing  interest  in  the  properties  of  the  elements 
themselves  focused  the  attention  of  chemists  about  1860 
upon  the  relation  of  properties  to  atomic  weights.  Several 
conditions  fostered  this  principle  of  classification.  One  was 
the  atomic  weight  determinations  of  Stas,  whose  masterly 
work  yielded  very  exact  atomic  weights.  Another  was  the 
acceptance  by  most  chemists  of  a  uniform  table  of  atomic 
weights.  A  third  was  the  vast  accumulation  of  facts  about 
the  elements  and  their  compounds.  Chemists  were  ready 
for  a  new  and  broader  classification  of  the  elements. 

The  Periodic  Classification  of  the  Elements.  —  Previous 
to  1869  no  classification  included  all  the  elements.  In  that 
year  the  Russian  chemist  Mendelejeff  published  an  arrange- 
ment known  as  the  periodic  classification  of  the  elements. 
This  classification  revealed  a  new  relation  between  the 
properties  of  the  elements  and  their  atomic  weights,  and  is 
very  helpful  in  studying  the  chemical  elements.  Men- 
delejefTs  scheme,  slightly  modified  to  conform  to  subsequent 
discoveries,  is  substantially  as  follows  :  — 

If  all  the  elements  are  arranged  in  the  order  of  their  in- 
creasing atomic  weights,  a  series  results  in  which  similar 
or  closely  related  elements  occur  at  regular  intervals.  That 
is,  the  series  breaks  up  into  several  periods,  and  hence  the 
system  of  classification  is  called  periodic.  If  the  series  is 
divided  into  these  periods  and  the  periods  placed  below 
each  other,  a  table  is  secured  in  which  the  perpendicular 
columns  consist  of  groups  of  similar  elements,  i.e.  elements 
which  have  similar  relations  and  form  analogous  compounds 
which  have  similar  properties.  Such  a  table  is  shown  on 
the  following  page.  In  each  group  the  elements  have  been 
subdivided  into  families  in  order  to  emphasize  the  mutual 
relationship  among  closely  allied  elements. 


358 


INORGANIC  CHEMISTRY 


5 

*   o   s-. 

t—     OS     t- 

-     <5i    0 

°!  i-(  ««. 

S  S  S 

C>    ^    0 

35     CS     «> 

u 

g 

O 

II   II   II 

O    "^      C 

ii  ii  7 

lie 

h 

O 

0 
CO 

§ 

3 

S 

01 

1 

M 

II 

II 

g 

II 

^ 

c 

ta 

O 

^' 

« 

« 

0. 

t= 

S 

o 
o 

0 

CS 

S 

S 

0 
cS 

S 

i 

7 

1* 

II 

II 

II 

0 

II 

II 

II 

° 

0 

00 

o 

CO 

a 

H 

^ 

t1 

g 

o 
<* 

S3 

o 
»o 

t— 

0 

CO 

CS 

S 

iq 

03 

0 

o 

7 

PH 

J[ 

02 

6 

CO 

H 

S 

£ 

§ 

0, 

i 

c 

0 

7 

CO 
S 

co' 
H 

4) 

O 

o 

8 

CO 

c 
co 

IS 

(M 

.p 

i 

H 

e- 

I 
O 

0 

CQ 

| 

T-l 

^ 

co 

CS 

S 

cj 

o 

8j 
h 

CO 

jl 

c 

o 

7 
3 

0 

i 

H 

L 

S 

o 

^ 

o 

a. 

TH 

o 

CO 

6 

o5 

g 

§ 

« 

<M 

i 

o  . 

o 
1 

Cl 

5 

c 
S3 

1 

CO 

II 

8 

c3 

£f 

| 

M 

00 

CO 

§ 

o 

g 

^ 

So 

uD 

^ 

C5 

CO 

'"I 

CO* 

iO 

0 

CO 

Cs 

o 

^ 

<o 

co 

** 

1 

*J 

•r-l 

|| 

O 

s5 

a 

si 

II 

M 

s 
O 

^5 

PH 

tc 

O 

E 

0 

a. 

0, 

88 

8 

0 

a 

7 

II 

1 

7 

f 

0 

a 

£ 

4 

j 

H 

^ 

e 
o 

s 

H 

N 

CO 

4 

IQ 

0 

^ 

X 

05 

o 

H 

PH 

CLASSIFICATION   OF   THE   ELEMENTS  359 

From  the  table  it  is  seen  that  the  elements  fall  naturally 
into  two  large  general  classes  called  groups  and  periods. 
Those  elements  in  the  same  vertical  column  belong  to  the 
same  natural  group,  while  those  in  the  same  horizontal  row 
belong  to  the  same  period.  Selecting  from  each  group  the 
important  elements,  we  have  the  following  distinct  families: 
Group  0.  Inert  elements  or  argon  family  —  helium, 

neon,  argon,  krypton,  xenon,  niton. 
Group  I.     Alkali    metals    or    sodium    family  —  lithium, 

sodium,  potassium. 
Univalent  heavy  metals  or  copper  family  — 

copper,  silver,  gold. 
Group  II.     Alkaline  earth  metals  or  calcium  family  — 

calcium,  strontium,  barium,  radium. 
Bivalent  heavy  metals  or  zinc  family  —  mag- 
nesium, zinc,  cadmium,  mercury. 
Group  III.     Boron  family  —  boron. 

•  Earth    metals    or    aluminium    family  —  alu- 
minium. 
Group  IV.     Tetravalent  non-metals  or  carbon  family  — 

carbon,  silicon. 

Tetravalent  metals  or  tin  family  —  tin,  lead. 
Group  V.     Pentavalent  non-metals  and  metals  or  nitrogen 
family  —  nitrogen,  phosphorus,  arsenic,  anti- 
mony, bismuth. 

Group  VI.     Hexavalent   metals   or   chromium    family  — 

chromium,  molybdenum,  tungsten,  uranium. 

Hexavalent  non-metals   or  oxygen  family  — 

oxygen,  sulphur,  selenium,  tellurium. 
Group  VII.     Manganese  family  —  manganese. 

Halogen  elements  or  chlorine  family  —  fluo- 
rine, chlorine,  bromine,  iodine. 
Group  VIII.     Iron  family  —  iron,  cobalt,  nickel. 
Platinum  family  —  platinum. 


360  INORGANIC   CHEMISTRY 

In  several  of  these  families  the  similarity  is  marked, 
especially  between  the  alkali  metals  and  also  the  halogens  ; 
in  some  cases  the  resemblance  is  not  very  striking  except 
in  a  limited  number  of  properties.  So  far,  our  treatment 
has  ignored  these  groups,  brief  references  only  having  been 
made  to  the  argon  family.  In  the  succeeding  pages,  how- 
ever, the  different  groups  will  be  emphasized  both  as  to 
their  individual  members  and  their  collective  properties. 

As  stated  above,  the  elements  in  the  same  horizontal  row 
belong  to  the  same  period.  The  periodic  variations  of  the 
properties  of  certain  typical  elements  may  be  illustrated  by 
the  first  (I)  and  second  (II)  periods.  Ignoring  the  argon 
family  which  is  somewhat  anomalous,  and  beginning  with 
lithium,  the  general  chemical  properties  vary  regularly  with 
increasing  atomic  weight.  The  metallic  character  typified  by 
lithium  gradually  diminishes  through  beryllium  and  boron; 
while  the  feeble  non-metallic  character  typified  by  carbon 
increases  through  nitrogen  and  oxygen  until  fluorine  is 
passed  and  sodium  is  reached  ;  here  the  metallic  character 
reappears.  Proceeding  onward  from  sodium,  the  same 
decrease  of  basic  and  increase  of  acid  properties  is  noticed 
until  potassium  is  reached,  and  here  again  the  marked 
metallic  character  reappears.  There  is  no  sudden  change 
in  properties  until  we  pass  from  one  period  to  the  next. 
Thus,  fluorine  at  the  end  of  the  second  period  forms  a 
strong  acid,  but  sodium  at  the  beginning  of  the  third  period 
forms  a  strong  base.  Similarly,  chlorine  is  strongly  acidic, 
but  potassium,  which  is  the  first  metal  in  the  next  period, 
is  markedly  basic  ;  chlorine  is  a  typical  non-metal,  while 
potassium  is  a  typical  metal.  Not  all  the  periods  are  as 
typical  as  those  just  cited,  but  in  many  cases  the  progres- 
sive change  in  chemical  properties  is  too  obvious  to  ascribe 
to  mere  chance.  Many  of  these  relations  will  be  pointed  out 
in  the  following  chapters.  Indeed  it  is  hardly  possible  to 


CLASSIFICATION   OF   THE   ELEMENTS  361 

appreciate  the  full  significance  of  the  periodic  classification 
until  most  of  the  elements  have  been  studied. 

Periodic  Law.  —  The  relation  between  properties  and 
atomic  weights  which  brings  about  this  periodic  variation  is 
general,  and  is  often  summarized  in  the  form  of  a  law  known 
as  the  periodic  law,  thus  :  — 

The  properties  of  the  elements  are  periodic  functions  of  their 
atomic  weights. 

Function  here  means  the  exhibition  of  some  special  rela- 
tion, viz.  that  of  properties  to  atomic  weight.  Interpreted 
freely,  the  law  means:  (1)  properties  and  atomic  weight  are 
related,  they  depend  upon  each  other ;  and  (2)  this  relation 
is  exhibited  repeatedly  at  regular  intervals. 

Conclusion.  —  An  examination  of  the  table  on  page  358 
shows  some  imperfections  in  the  periodic  classification. 
For  example,  there  are  gaps.  These  probably  correspond 
to  elements  not  yet  discovered  or  fully  investigated.  Three 
such  gaps,  which  were  in  the  original  table,  have  been  filled. 
When  Mendelejeff  proposed  his  arrangement,  he  predicted 
the  discovery  of  three  elements  having  definite  properties. 
These  elements  —  gallium,  scandium,  and  germanium  — 
have  since  been  discovered,  and  now  occupy  their  predicted 
places  in  the  table.  Possibly  other  gaps  will  be  filled  by 
newly  discovered  elements.  Several  elements  do  not  fall 
into  their  proper  places.  Thus,  the  atomic  weight  of  argon 
indicates  that  this  element  should  exchange  places  with 
potassium.  Similarly  the  positions  of  iodine  and  tellurium 
should  be  reversed ;  their  properties,  however,  necessitate 
the  present  places.  Hydrogen  has  no  appropriate  place  in 
the  series. 

The  periodic  classification,  although  imperfect  in  some 
particulars,  simplifies  the  study  of  chemistry,  and  will  be 
utilized  in  its  larger  aspects  in  many  of  the  following  chapters. 


CHAPTER  XX 
Fluorine  —  Bromine  —  Iodine 

FLUORINE,  bromine,  and  iodine,  together  with  chlorine, 
constitute  a  family  in  the  seventh  (VII)  periodic  group, 
often  called  the  halogens.  They  have  similar  chemical  prop- 
erties, and  form  analogous  compounds  which  likewise  have 
similar  properties,  differing  mainly  in  degree. 

Halogen  means  "  sea-salt  producer."  It  is  applied  to  this 
group  of  elements  because  they  form  salts  which  are  found  in 
sea  water  and  resemble  sodium  chloride  (common  salt  or  sea 
salt).  Chlorides,  bromides,  fluorides,  and  iodides  are  called 
halides.  The  Greek  word  for  salt,  hals,  suggested  these 
descriptive  terms. 

Chlorine  was  fully  discussed  in  Chapter  XII. 

FLUORINE 

Occurrence.  —  Fluorine  is  the  most  active  of  all  the  halogen 
elements,  and  is  never  found  free  in  nature.  It  occurs 
abundantly  in  combination  with  calcium  as  fluor  spar  or 
calcium  fluoride  (CaF2).  Other  native  compounds  are 
cryolite  (Na3AlF6)  and  apatite  (Ca5F(PO4)3).  Minute  quan- 
tities of  combined  fluorine  are  found  in  bones  and  blood, 
in  the  enamel  of  the  teeth,  and  in  sea  and  some  mineral 
waters. 

The  Isolation  of  Fluorine  was  accomplished  in  1886  by 
Moissan,  though  many  unsuccessful  attempts  had  previously 
been  made.  He  decomposed  hydrofluoric  acid  by  electricity, 
and  collected  the  liberated  fluorine.  The  achievement  was 

362 


MOISSAN 
1852-1907 


FLUORINE 


363 


attended  with  tremendous  difficulties,  owing  to  the  intense 
activity  of  the  fluorine  and  its  corrosive  properties. 

The  essential  parts  of  the  apparatus  used  by  Moissan  are  shown 
in  Figure  56.  The  U-tube,  made  of  an  alloy  of  platinum  and  iridium, 
is  provided  with  tightly  fitting  stoppers  of 
fluor  spar  (S,  S).  Through  the  stoppers  pass 
the  electrodes  (E,  £"),  of  platinum-iridium, 
held  in  place  by  screw  caps  (C,  C).  Side 
tubes  (7\  71)  allow  the  liberated  gases  (fluorine 
and  hydrogen)  to  be  drawn  off  separately 
through  platinum  delivery  tubes.  Perfectly 
dry  hydrofluoric  acid  is  put  into  the  U-tube, 
and  dry  acid  potassium  fluoride  (HKF2)  is 
added  to  enable  the  solution  to  conduct  the 
current  —  liquefied  hydrofluoric  acid  itself 
being  a  non-conductor.  The  U-tube  is  then 
cooled  to  a  low  temperature  (—  23  to  —  50°  C.), 
and  on  passing  a  current  through  the  solution 

fluorine  is  evolved  at  the  positive  electrode    JPIG.  55. Moissan 'sap- 

(anode)  and   hydrogen  at  the  negative  elec-       paratus  for  preparing 
trode    (cathode).      The   fluorine  freed   from       fluorine, 
hydrofluoric    acid   vapor    was    collected    by 

Moissan  at  first  in  a  platinum  tube  with  a  thin  fluor  spar  plate  closing 
each  end,  so  he  could  look  inside  and  examine  the  gas.  Later  he 
found  that  the  electrolysis  could  be  performed  in  a  copper  U-tube 
and  pure  dry  fluorine  could  be  collected  in  a  glass  tube. 

Properties.  —  Fluorine  is  a  gas  having  a  greenish  yellow 
color,  though  lighter  and  more  yellowish  than  chlorine. 
The  critical  temperature  is  —  120°  C.  Subjected  to  pressure 
and  a  sufficiently  low  temperature,  it  condenses  to  a  very 
pale  yellow  liquid,  which  boils  at  about  —  187°  C.  The 
vapor  density  shows  that  the  molecular  weight  is  about  38 ; 
hence,  each  molecule  contains  two  atoms  and  fluorine  gas 
has  the  formula  F2  (the  atomic  weight  being  19). 

Chemically  fluorine  is  intensely  active.  Hydrogen,  bro- 
mine, iodine,  sulphur,  phosphorus,  carbon,  silicon,  and  boron 


364  INORGANIC   CHEMISTRY 

take  fire  in  it,  but  oxygen,  nitrogen,  and  argon  do  not  unite 
with  it.  Most  metals  burn  in  it,  forming  fluorides.  Gold 
and  platinum  are  not  attacked  by  it  below  red  heat.  Copper 
becomes  coated  with  copper  fluoride,  which  protects  the 
metal,  so  that  copper  vessels  may  be  used  as  fluorine  gen- 
erators and  reservoirs.  Moissan  used  a  copper  U-tube  to 
prepare  large  volumes.  Water  is  decomposed  by  it  at 
ordinary  temperatures,  while  hydrocarbons  are  instantly 
decomposed,  hydrofluoric  acid  and  carbon  fluorides  being 
the  important  products. 

Hydrofluoric  Acid  is  the  compound  of  fluorine  correspond- 
ing to  hydrochloric  acid.  It  is  prepared  by  the  interaction  of 
a  fluoride  and  concentrated  sulphuric  acid,  just  as  hydro- 
chloric acid  is  prepared  from  a  chloride  and  sulphuric  acid. 
Calcium  fluoride  is  usually  used,  and  the  experiment  is  per- 
formed preferably  in  a  lead  dish.  The  chemical  change  is 
represented  thus  :  — 


CaF2       +       H2S04       =        2HF         +        CaSO4 

Calcium  Sulphuric  Hydrofluoric  Calcium 

Fluoride  Acid  Acid  Sulphate 

Pure  hydrofluoric  acid  is  a  colorless  liquid  which  boils  at 
about  19°  C.  It  is  very  volatile  and  is  readily  transformed 
into  a  colorless  gas  (hydrogen  fluoride),  which  fumes  in  the 
air  and  dissolves  in  water,  the  aqueous  solution  being  the 
commercial  hydrofluoric  acid.  Both  gas  and  solution  are 
dangerous  substances.  The  gas  is  extremely  poisonous,  and 
the  liquid,  if  dropped  on  the  skin,  produces  terrible  sores. 
Owing  to  its  corrosive  action  hydrofluoric  acid  is  preserved 
and  sold  in  hard  rubber,  lead,  or  wax  bottles.  The  vapor 
density  of  hydrofluoric  acid  gas  varies  with  the  temperature. 
From  19°  to  30°  C.  the  vapor  density  indicates  a  molecular 
weight  of  40,  while  at  88°  C.  the  molecular  weight  is  20. 
The  atomic  weight  of  fluorine  is  19.  Hence  at  the  lower 


BROMINE  365 

temperature  the  formula  of  the  gas  is  H2F2  and  at  the  higher 
temperature  it  is  HF. 

Hydrofluoric  acid  interacts  with  metals,  thereby  liberating 
hydrogen  and  forming  fluorides.  It  also  interacts  with'bases 
and  metallic  oxides  and  forms  fluorides.  There  are  two 
classes  of  fluorides  —  normal  and  acid.  Thus,  calcium 
fluoride  (CaF2)  is  a  normal  salt  and  hydrogen  potassium 
fluoride  (HKF2)  is  an  acid  salt. 

A  solution  of  hydrofluoric  acid  and  the  moist  gas  both 
attack  glass  and  are  used  extensively  in  etching.  The  glass 
is  coated  with  wax,  and  the  design  to  be  etched  is  scratched 
through  the  wax.  The  glass  is  then  exposed  to  the  gas  or 
liquid,  which  attacks  the  unprotected  places.  When  the 
wax  is  removed,  a  permanent  etching  like  the  design  is  left. 
Glass  is  an  artificial  compound  of  silicon  —  a  silicate.  The 
corrosive  action  of  hydrofluoric  acid  upon  glass  is  due  to  the 
ease  with  which  the  acid  decomposes  glass  and  forms  with 
silicon  a  volatile  compound,  called  silicon  tetrafluoride  (SiF4). 
Equations  for  the  essential  chemical  changes  are :  — 

CaSi03     +     6HF     =     SiF4     +     CaF2     +     3  H2O 

Calcium  Hydrofluoric          Silicon  Calcium  Water 

Silicate  Acid  Tetrafluoride       Fluoride 

SiO2        +         4  HF  SiF4        +         2  H2O 

Silicon  Hydrofluoric  Silicon 

Dioxide  Acid  Tetrafluoride 

Scales  on  thermometers  and  on  other  graduated  glass  instru- 
ments are  often  etched  with  hydrofluoric  acid. 

BROMINE 

Occurrence.  —  Bromine  is  never  found  free  in  nature  on 
account  of  its  chemical  activity.  Bromides  are  widely  dis- 
tributed, especially  sodium  bromide  and  magnesium  bromide. 
The  salt  springs  of  Ohio,  West  Virginia,  Pennsylvania,  and 


366  INORGANIC  CHEMISTRY 

Michigan,  and  the  salt  deposits  at  Stass- 
furt  in  Germany  furnish  the  main  supply 
of  the  element.  Sea  water,  Chile  saltpeter 
(NaNO3),  and  certain  seaweeds  contain  a 
small  quantity  of  combined  bromine. 

Preparation.  —  Bromine  is  obtained  from 
its  native  compounds   by   electrolysis   and 
by    treatment    with    chlorine  or   with  sul- 
phuric acid  and  some  oxidizing  agent  like 
FIG.  57.  —  Appara-  potassium  chlorate  or    manganese  dioxide. 
brom°inePrir1he  In  the  laboratory>  bromine  is  prepared  by 
laboratory.  Part  heating  potassium  bromide  with  manganese 
of  the  liberated  dioxide  and  sulphuric  acid  in  a  glass  ves- 


part  quickly  con-  ated  as  a  dense,  brown  vapor,  which  often 
densestoaliquid  condenses  to  a  liquid  and  runs  down  the 

rh^^bendoahe  walls  of  the  vesseL     The  chemical  change 
delivery  tube.       is  represented  thus  :  — 

2  KBr  +  2  H2S04  +  Mn02  =  Br2  +  MnSO4  +  K£04  +  2  H2O 


Potas-         Sulphuric        Manga-       Bro-         Manga-          Potas-          Water 
sium  Acid  nese          mine  nese  sium 

Bromide  Dioxide  Sulphate       Sulphate 

Bromine  is  sometimes  prepared  in  the  laboratory  by  treating 
a  bromide  with  manganese  dioxide  and  hydrochloric  acid. 

Bromine  is  obtained  technically  from  the  solution  left  after 
sodium  chloride  is  crystallized  from  brine  and  also  from  the  liquor 
left  after  potassium  chloride  is  extracted  from  impure  carnallite 
(KC1,  MgCl2  .  6  H2O).  The  hot  liquid  flows  down  a  large  tower 
filled  with  clay  balls  ;  chlorine  gas  (frequently  obtained  by  the 
electrolysis  of  chlorides)  and  steam  forced  in  at  the  bottom  meet 
the  liquid  and  liberate  the  bromine,  which  passes  as  a  vapor  out 
of  the  top  into  a  condenser.  Equations  for  the  changes  are  :  — 

MgBr2  +         C12       =      Br2         +  MgCl2 

Magnesium  Bromide  Chlorine  Bromine  Magnesium  Chloride 

2NaBr  +         C12       =      Br2         +         2  NaCl 

Sodium  Bromide  Chlorine  Bromine  Sodium  Chloride 


BROMINE  36? 

In  the  periodic  process,  used  chiefly  in  the  United  States,  a  huge 
stone  still  is  charged  with  manganese  dioxide,  hot  bittern,  and  sul- 
phuric acid,  and  heated  by  steam.  The  bromine  distills  into  a  con- 
denser, as  in  the  other  process.  Sometimes  potassium  chlorate  is 
used  as  the  oxidizing  agent,  and  the  equation  for  the  essential 
chemical  change  is  :  — 

KC1O3       +       6HC1       =       3C12      +      KC1      +     3  H2O 

Potassium  Hydrochloric  Chlorine  Potassium  Water 

Chlorate  Acid  Chloride 

Properties.  —  Bromine  is  a  heavy,  reddish  brown  liquid  at 
the  ordinary  temperature.  Its  specific  gravity  is  about  3. 
It  is  a  volatile  liquid,  boiling  at  about  59°  C.  The  vapor, 
which  is  given  off  freely,  has  a  disagreeable,  suffocating  odor. 
This  property  suggested  the  name  bromine  (from  the  Greek 
word  bromos,  a  stench).  It  is  poisonous,  and  burns  the  flesh 
frightfully.  Bromine  is  somewhat  soluble  in  water.  The 
solution  called  bromine  water  has  a  brown  color  and  is  some- 
times used  instead  of  bromine  itself.  The  solution  contains 
an  unstable  bromine  hydrate.  Bromine  dissolves  in  carbon 
disulphide,  and  this  solution  is  yellow.  Its  vapor  density 
up  to  about  750°  C.  requires  a  molecular  weight  of  160,  and 
since  the  atomic  weight  is  79.92;  the  formula  of  bromine 
vapor  is  Br2. 

Bromine  combines  directly  with  many  elements,  especially 
hydrogen,  phosphorus,  and  metals.  The  action  is  not  so 
violent  as  with  chlorine.  In  fact  free  chlorine  readily  dis- 
places bromine  from  some  of  its  compounds.  The  chemical 
properties  of  bromine  can  be  illustrated  by  a  simple  experi- 
ment. If  powdered  magnesium  is  added  to  bromine  water, 
the  brown  color  disappears,  owing  to  the  formation  of  color- 
less magnesium  bromide  by  the  direct  combination  of  mag- 
nesium and  bromine.  Upon  the  addition  of  chlorine  (or 
chlorine  water)  to  this  colorless  solution,  the  brown  color 
reappears,  owing  to  the  free  bromine  which  is  displaced  from 
the  magnesium  bromide  by  the  more  active  chlorine. 


368  INORGANIC  CHEMISTRY 

Compounds  of  Bromine.  —  Hydrogen  bromide  (HBr)  is 
a  colorless  pungent  gas,  which  fumes  in  the  air  and  dissolves 
freely  in  water,  forming  a  solution  usually  called  hydrobromic 
acid.  Its  other  properties  closely  resemble  those  of  hydro- 
chloric acid,  though  it  is  less  stable  than  its  chlorine  ana- 
logue. Bromides  are  Salts  of  hydrobromic  acid,  though 
many  are  formed  by  direct  combination  with  bromine. 
Most  bromides,  like  the  chlorides,  dissolve  in  water.  Po- 
tassium bromide  (KBr)  is  a  white  solid,  made  by  decom- 
posing iron  bromide  with  potassium  carbonate.  It  is  used 
extensively  as  a  medicine  and  in  photography  (in  preparing 
silver  bromide  plates  and  films). 

Miscellaneous.  —  Bromine  itself  is  used  to  make  potassium 
bromide  and  other  compounds,  especially  certain  coal  tar 
dyes. 

Balard  discovered  bromine  in  1826  in  the  mother  liquor 
(or  bittern)  from  brine.  Liebig,  to  whom  it  was  submitted 
for  examination,  supposed  it  was  chloride  of  iodine,  and  thus 
failed  to  discover  its  elementary  nature,  because,  as  he  said, 
he  yielded  to  "  explanations  not  founded  on  experiment." 

IODINE 

Occurrence.  —  Free  iodine  is  never  found  in  nature,  but 
like  chlorine  and  bromine  it  occurs  in  combination  with 
metals,  especially  sodium,  potassium,  or  magnesium.  It  is 
widely  distributed,  though  the  quantity  in  any  one  place  is 
small.  Tobacco,  water-cress,  cod-liver  oil,  oysters,  and 
sponges  contain  minute  quantities  of  iodine  compounds. 
Native  iodides  of  silver  and  of  mercury  are  found.  The  ash 
of  some  seaweeds  contains  from  0.5  to  1.5  per  cent  of  its 
weight  of  iodides  of  sodium  and  potassium.  Sodium  iodate 
(NaIO8)  occurs  in  the  deposits  of  saltpeter  in  Chile,  and  is 
now  the  main  source  of  the  element. 


IODINE  369 

Preparation.  —  Iodine  is  prepared  in  the  laboratory  by  a 
method  similar  to  that  used  for  bromine.  Potassium  iodide, 
manganese  dioxide,  and  sulphuric  acid  are  heated  in  a  glass 
vessel,  and  the  iodine  is  liberated  as  a  violet  vapor,  which 
quickly  condenses  on  the  upper  part  of  the  vessel  as  dark 
grayish  crystals.  The  equation  for  the  chemical  change  is  — • 

2  KI  +  2  H2SO4  +  Mn02  =  I,  +  MnSO4  +  K2SO4  +  2  H2O 

Potassium  Sulphuric     Manganese  Iodine  Manganese      Potassium      Water 
Iodide         Acid  Dioxide  Sulphate         Sulphate 

On  a  commercial  scale  iodine  is  prepared  from  the  ash  of  seaweeds 
and  from  the  mother  liquors  of  Chile  saltpeter.  (1)  Along  the  coasts 
of  France,  Japan,  Scotland,  and  Norway  seaweed  is  collected  and 
burned;  sometimes  the  seaweed  is  allowed  to  ferment  before  being 
burned  in  order  to  convert  the  complex  iodine  compounds  into  non- 


FIG.  58.  —  Apparatus  for  purifying  iodine. 

volatile  iodides.  From  the  ash,  which  is  called  kelp  or  varec,  the 
soluble  portions  are  removed  by  agitation  with  water.  The  filtered 
liquid  is  further  purified,  and  from  the  final  mother  liquor  in  which 
the  iodides  are  dissolved  the  iodine  is  extracted  by  heating  with 
sulphuric  acid  and  manganese  dioxide.  Sometimes  chlorine  is  used 
to  extract  the  iodine,  the  equation  for  this  chemical  change  being  — • 

2  Nal     +      Cla      =     I2     +     2  NaCl 

Sodium  Chlorine       Iodine  Sodium 

Iodide  Chloride 

In  either  case  the  mother  liquor  and  its  added  ingredients  are  dis- 
tilled gently  in  an  iron  pot  with  a  lead  cover,  which  is  connected  with 
two  rows  of  bottle-shaped  condensers  (Fig.  58).  The  iodine,  which 


370  INORGANIC   CHEMISTRY 

collects  in  these  condensers,  is  purified  by  washing  and  resubliming. 
(2)  In.  the  other  process  the  mother  liquor  (which  contains  about 
22  per  cent  of  sodium  iodate)  from  the  Chili  saltpeter  is  mixed 
with  sodium  sulphite  and  acid  sodium  sulphite  (HNaSO3);  the 
precipitated  iodine  is  collected  on  coarse  cloth,  washed,  dried,  and 
then  resublimed,  as  described  above.  The  equation  for  this  method 
of  preparation  is  — 

2NaIO3    4-    3Na2SO3  +  2HNaSO3  =  I2  -f  5Na2SO4    +    H2O 

Sodium  Sodium  Acid  Sodium      Iodine         Sodium  Water 

Iodate  Sulphite  Sulphite  Sulphate 

Courtois,  a  French  chemist,  discovered  iodine  in  1812,  in  an  attempt 
to  prepare  potassium  nitrate  from  seaweed.  Davy  and  Gay-Lussac 
established  its  elementary  nature  and  discovered  many  of  its  proper- 
ties. The  present  name  was  given  by  Davy. 

Properties.  —  Iodine  is  a  dark  grayish  crystalline  solid, 
resembling  graphite  in  luster.  It  crystallizes  in  plates  which 
have  a  specific  gravity  of  about  5.  It  is  volatile  at  the  ordi- 
nary temperature,  melts  at  114°  C.,  and  boils  at  184°  C. 
When  gently  heated,  the  vapor  which  is  formed  has  a  beauti- 
ful violet  color.  This  color  suggested  the  name  of  iodine 
(from  the  Greek  word  iodes,  violetlike).  The  vapor  is  nearly 
nine  times  heavier  than  air.  It  has  an  odor  resembling  that 
of  dilute  chlorine,  though  less  irritating.  When  the  vapor 
is  heated,  its  color  changes  from  violet  to  deep  blue,  and  the 
density  decreases.  Experiment  shows  that  from  about 
200  to  700°  C.  the  vapor  density  requires  a  molecular  weight 
of  about  255;  since  the  atomic  weight  is  126.92,  the  mole- 
cules contain  only  two  atoms  and  the  formula  of  iodine 
vapor  is  I2  up  to  this  temperature.  As  the  temperature 
rises  the  molecules  dissociate,  until  at  about  1700°  C.,  the 
vapor  consists  entirely  of  atoms.  Iodine  stains  the  skin 
yellow,  and  turns  cold  starch  solution  blue.  The  presence 
of  a  minute  trace  of  iodine  may  be  thus  detected,  one  part 
of  iodine  in  over  400,000  parts  of  water  producing  the  blue 
color.  The  exact  nature  of  this  blue  substance  is  unknown. 


IODINE  371 

The  presence  of  starch  in  many  vegetable  substances  can  be 
shown  by  this  delicate  test.  Iodine  dissolves  slightly  in 
water,  and  freely  in  alcohol,  chloroform,  carbon  disulphide, 
and  potassium  iodide  solution.  The  chloroform  and  carbon 
disulphide  solutions  are  violet,  but  the  others  are  brown  or 
even  black. 

The  chemical  properties  of  iodine  resemble  those  of  chlorine 
and  bromine,  but  it  is  less  active.  It  forms  no  hydrate  with 
water,  differing  from  chlorine  and  bromine  in  this  respect. 
Bromine  and  chlorine  displace  iodine  from  its  binary  com- 
pounds, chlorine  and  chlorine  water  being  often  used  for  this 
purpose.  It  combines  directly  with  some  non-metals  and 
most  metals. 

Compounds  of  Iodine  resemble  the  corresponding  ones  of 
chlorine  and  bromine.  Hydriodic  acid  (HI)  is  much  like 
hydrobromic  and  hydrochloric.  Iodides  are  salts  of  hy- 
driodic  acid,  and  like  many  salts  they  are  prepared  in  various 
ways.  In  general  behavior  they  are  similar  to  bromides  and 
chlorides.  Potassium  iodide  (KI)  is  prepared  and  used  like 
potassium  bromide. 

Brief  reference  was  made  in  Chapter  XI  to  the  equilibrium 
established  when  hydriodic  acid  is  formed  by  heating  hydro- 
gen and  iodine  in  a  closed  tube.  The  elements  begin  to 
combine  at  about  445°  C.,  while  the  acid  begins  to  decom- 
pose at  about  180°  C.  The  reactions  proceed  in  direct 
and  reverse  directions,  though  not  to  completion,  for  a  con- 
dition of  equilibrium  is  soon'  reached  and  further  accumula- 
tion of  the  products  ceases.  Chemical  change  is  still  going 
on,  but  one  reaction  neutralizes  the  other.  For  example,  at 
about  445°  C.  a  condition  of  equilibrium  prevails,  both  reac- 
tions are  still  in  operation,  though  the  tube  contains  about 
79  per  cent  of  hydriodic  acid  and  21  per  cent  of  hydrogen 
and  iodine.  This  state  of  equilibrium  is  maintained  as  long 


372  INORGANIC   CHEMISTRY 

as  certain  conditions  prevail,  but  like  other  states  of  equilib- 
rium it  is  rather  sensitive  and  is  easily  displaced  by  varying 
the  temperature,  the  pressure,  or  the  concentration  of  the 
substances  in  the  tube.  The  equation  expressing  the  revers- 
ible reaction  just  discussed  may  be  written  — 


Miscellaneous.  —  Iodine  dissolved  in  alcohol,  in  potassium 
iodide  solution,  or  both,  is  used  (under  the  name  of  tincture 
of  iodine)  as  an  application  for  the  skin  to  prevent  the  spread 
of  eruptions  or  to  reduce  swellings.  Iodine  is  used  to  make 
medicinal  preparations,  especially  iodoform  (CHI3),  which 
is  used  as  an  antiseptic  dressing  for  wounds.  Large  quan- 
tities of  iodine  are  used  in  making  aniline  dyes. 

Halogen  Family.  —  The  physical  and  chemical  properties 
of  the  halogen  elements  furnish  a  typical  illustration  of  the 
resemblances,  differences,  and  gradation  of  properties  which 
characterize  a  family  of  elements  in  the  same  periodic  group- 
If  these  elements  are  arranged  in  the  order  of  their  atomic 
weights,  from  fluorine  (19.0)  through  chlorine  (35.46)  and 
bromine  (79.92)  to  iodine  (126.92),  the  periodic  nature  of 
the  group  is  revealed.  Thus,  the  specific  gravity  increases 
in  this  order  (i.e.  fluorine  to  iodine)  ;  the  color  likewise  grows 
deeper,  but  the  volatility  decreases.  So  also,  the  melting 
points  of  the  solidified  elements  and  the  boiling  points  of 
the  liquefied  elements  increase  in  this  order.  Chemically, 
these  elements  unite  with  hydrogen  to  form  analogous  com- 
pounds, and  the  intensity  of  the  chemical  action  decreases 
gradually  as  we  pass  from  fluorine  to  iodine. 

The  halogen  acids  resemble  each  other  in  their  solubility, 
all  (except  HF)  forming  solutions  in  which  the  dissociation 
is  large;  in  other  words,  they  are  strong  acids.  Salts  of  the 


IODINE  373 

halogen  acids  often  resemble  one  another,  especially  those 
of  sodium  and  potassium,  which  are  white  solids,  soluble  in 
water,  from  which  they  crystallize  in  cubes. 

The  valence  of  the  halogens  is  one  toward  hydrogen  and 
metals. 

PROBLEMS  AND  EXERCISES 

1.  What  is  the  percentage  composition  of   (a)  fluor  spar  (CaF2) 
and  (6)  cryolite  (Na3AlF6)? 

2.  How  much  (a)  calcium  sulphate  and  (6)  hydrofluoric  acid  are 
formed  by  heating  100  gm.  of  fluor  spar  with  sulphuric  acid? 

3.  Calculate  the  percentage  composition  of  (a)  potassium  bromide 
(KBr),   (6)  potassium  iodide  (KI),   (c)  silver  bromide  (AgBr),  and 
(d)  iodoform  (CHI3). 

4.  How  much  potassium  iodide  is  needed  to  prepare  63.5  gm.  of 
iodine  ? 

5.  How  much  potassium  bromide  is  needed  to  prepare  10  gm.  of 
bromine  ? 

6.  If  3.946  gm.  of  silver  are  needed  to  precipitate  the  bromine  in 
4.353  gm.  of  potassium  bromide,  what  is  the  atomic  weight  of  bro- 
mine? 

7.  If  the  specific  gravity  of  bromine  is  3,  what  volume  does  one 
pound  occupy? 

8.  If  63.8351  gm.  of  silver  iodide  yield  38.9656  gm.  of  silver  chloride, 
what  is  the  atomic  weight  of  iodine  if  107.88  and  35.46  are  accepted 
as  the  atomic  weights  of  silver  and  chlorine  respectively? 

9.  Write  the  formulas  of  the  following  compounds  by  applying 
the  principle  of  valence  (see  Chapter  XIV) :    Lithium  chloride,  mag- 
nesium iodide,  mercurous  iodide,  mercuric  bromide,  aurous  chloride, 
auric  bromide,  barium  fluoride,  silicon  fluoride,  platinum  chloride, 
zinc  chloride,  ferrous  bromide,  ferric  chloride,  lead  iodide. 

10.  Write  the  equations  for  the  following  reactions:   (a)  Potassium 
bromide  and  silver  nitrate  form  silver  bromide  and  potassium  nitrate. 
(6)  Sodium  iodide  and  silver  nitrate  form  silver  iodide  and  sodium 
nitrate.     Write  the  corresponding  ionic  equations. 

11.  Calculate  the  simplest  formulas  :  (a)  F  =  48.72,  Ca  =  51.28 ; 
(6)  Br  =  67.22,  K  =  32.77;  (c)   I  =  76.5,  K  =  23.49. 

12.  The  formulas  of  bromic  acid  and  iodic  acid  are  HBrOs  and 
HIO3  respectively.     Write  the  formulas  of  their  salts  corresponding 
to  K,  calcium,  Mg,  sodium,  and  Ba. 


CHAPTER  XXI 
Boron 

Occurrence.  —  Boron  (B)  is  never  found  free,  but  several 
of  its  compounds  are  abundant,  e.g.  borax  (Na2B4O7),  boric 
acid  (H8BO8),  boracite  ((MgaBgO^j .  MgCl2),  and  colemanite 
(Ca2B6Ou.5H2O). 

Boron  is  prepared  by  heating  the  oxide  (B2O3)  with  mag- 
nesium or  the  chloride  (BC13)  with  hydrogen.  It  is  black, 
non-crystalline,  and  very  hard.  It  unites  with  oxygen  and 
with  nitrogen,  forming  the  oxide  (B2O3)  and  a  nitride  (BN). 
When  heated  in  the  electric  furnace  with  carbon,  it  forms 
carbon  boride  (CB6),  which  is  nearly  as  hard  as  diamond. 

Boric  Acid,  H8BO8,  is  contained  in  the  waters  and  steam  of 
certain  volcanic  regions,  notably  Tuscany  in  Italy.  Large 
basins  or  tanks  are  built  around  these  steam  jets  (called 
suffioni)  and  are  arranged  so  that  the  water  flows  at  intervals 
from  one  reservoir  into  the  next  lower,  constantly  becoming 
charged  with  more  boric  acid,  as  the  steam  condenses.  The 
final  solution  is  evaporated  by  aid  of  the  heat  from  the  steam 
jets,  and  the  crude  boric  acid  which  settles  out  is  purified 
by  recrystallization.  This  compound  is  sometimes  called 
boracic  acid.  Boric  acid  is  made  in  the  United  States  from 
borax  and  colemanite,  and  in  Germany  from  the  boracite 
found  at  Stassfurt.  The  essential  feature  of  the  process  is  a 
decomposition  of  the  mineral  into  boric  acid  by  an  inorganic 
acid.  The  following  equation  illustrates  the  transforma- 
tion: — 

Na,B4OT  +  Hj.804  +  5H,O    =    4H8BO8     +     Na2SO4 

Borax  Sulphuric  Boric  Sodium 

Acid  Acid  Sulphate 

374 


BORON  375 

Boric,  acid  crystallizes  in  lustrous,  white  flakes,  which  fee) 
greasy.  It  dissolves  slightly  in  cold  wafer,  readily  in  liot 
wafer  :ind  in  alcohol.  When  I  lie  Alcoholic  solution  is  burned, 
;i  l)oron  compound  colors  tlio  vapor  green.  This  is  a  test 
for  boron  compounds. 

I'oric  arid  is  used  in  making  borax,  in  the  manufacture  of 
enamels  .-Hid  gla/,es  for  pottery,  .MS  :in  ;i  nliseptic.  in  medicine 
and  snidery,  and  for  preserving  meat,  lisli,  milk,  butter,  beer, 
.MIX!  wine. 

When  boric,  acid  is  heated,  it  loses  water  and  is  transformed 
into  metaboric  acid  (HBOa)  at  100°  C.  and  tctraboric  acid 
(Ilal^O;)  Jit  J4()°C.  Boric  acid  forms  no  salts,  but.  the  other 
acids  do,  the  best-known  salt  being  sodium  tctraborato  or 
borax. 

Borax,  Na2B4Or,  occurs  native  in  California,  and  an  impure 
borax  called  tinkal  comes  from  the  Kast.  Most  of  the  com- 
mercial borax  is  made  from  calcium  borate  (colemanite)  by 

boiling  with  sodium  carbonate  thus: 

2  Ca2BflOn  +  4  Na2C08+H2O  =  3  Na2B407+4  CaC08+2  NaOH 

Calcium  Sodium  Borax  Calcium        Hodiniu 

Ilydroxido 


Borax  is  a  white  crystalline  solid  and  has  five  or  ten  mole- 
cules of  water  of  crystallization,  depending  upon  the  tem- 
perature at  which  crystalli/;il  ion  occurs.  It  effloresces  in 


Fio.  59.  —  Looped  platinum  wire  for  m.ikiriK  towts  with  borax  beads. 

the  air.  When  crystallized  borax  is  heated,  it  swells  up  (as 
the  water  of  crystallization  escapes)  into  a  white  porous 
mass,  which  finally  becomes  a  glassy  solid.  If  the  borax 
is  melted  on  the  end  of  a  looped  platinum  wire,  the  trans- 
parent globule  is  called  a  borax  bead  (Fig,  59),  This 


376 


INORGANIC  CHEMISTRY 


glassy  borax  dissolves  metallic  substances,  especially  metallic 
oxides.  These  beads  may  be  made  to  assume  different  colors 
characteristic  of  the  metals,  when  the  beads  are  heated  with 
the  oxides  or  solutions  of  different  metals.  The  -following 
table  shows  the  — 

COLORS  OF  BORAX  BEADS 


OXIDIZING  FLAMK 

REDUCING  FLAME 

METAL 

Hot 

Cold 

Hot 

Cold 

Chromium 

Reddish  yellow 

Yellowish  green 

Green 

Green 

Cobalt 

Blue 

Blue 

Blue 

Blue 

Copper 

Green 

Greenish  blue 

Colorless 

Red 

Manganese 

Violet 

Violet 

Colorless 

Colorless 

The  bead  test  is  often  used  in  chemistry  to  confirm  other 
observations  or  to  suggest  further  examination.  The  chem- 
ical changes  in  borax  beads  can  be  readily  understood  if 
borax  is  regarded  as  a  mixture  of  sodium  metaborate  and 
boron  trioxide  (2  NaB02  -f-  B2O3) ;  the  acid  oxide  (B2O8) 
unites  with  the  basic  (i.e.  metallic)  oxides  and  forms  a  colored 
borate.  For  example,  with  copper  oxide  a  green  bead  is 
obtained  which  is  2NaB02.  Cu(BO2)2. 

A  solution  of  borax  has  an  alkaline  reaction,  because  borax 
hydrolyzes  with  water  and  boric  acid  is  only  slightly  ionized. 

Large  quantities  of  borax  are  used  in  the  manufacture  of 
enamels  and  glazes,  especially  those  which  form  the  protective 
coating  of  domestic  utensils.  Considerable  is  used  for 
preserving  canned  meat  and  fish.  It  is  a  cleansing  agent, 
and  large  quantities  are  consumed  in  laundries  as  well  as  in 
the  manufacture  of  soap,  particularly  the  variety  intended 
for  use  in  hard  water.  (See  Soap  and  Hard  Water.)  The 
property  of  dissolving  oxides  adapts  borax  for  use  in  soldering 
certain  metals.  Solder  adheres  only  to  clean  metals,  so  a 


BORON  377 

little  borax  is  used  to  dissolve  the  film  of  oxide  on  the  surfaces 
to  be  joined.  It  also  finds  use  as  a  mordant  in  calico  printing 
and  in  dyeing,  and  in  the  manufacture  of  water-soluble 
varnishes.  It  is  an  ingredient  of  the  ointments,  lotions,  and 
powders  which  are  designed  to  relieve  hoarseness  or  skin 
eruption. 

Miscellaneous.  —  Boron  is  a  non-metal  and  belongs  to  the 
aluminium  family  in  the  third  periodic  group  (III),  but  chem- 
ically it  closely  resembles  carbon  and  silicon  and  their  com- 
pounds. Its  valence  is  three. 

PROBLEMS 

1.  What  per  cent  of  borax  (Na2B4O7. 10  H2O)  is  boron? 

2.  Write  the  equations  for  the  transformation  of  boric  acid  into 
metaboric  and  tetraboric  acids;    and  the  equation  for  the  formation 
of  boron  trioxide  from  tetraboric  acid. 

3.  How  many  grams  of  HaBOs  are  formed  by  the  interaction  of 
water,  sulphuric  acid,  and  782  gm.  of  borax? 

4.  Calculate  the  per  cent  of  boron  in  (a)  colemanite,  and   (6)  bo- 
racite. 

5.  How  many  grams  of  borax  can  be  made  from  a  metric  ton  of 
colemanite  ? 

6.  Write  the  formulas  of  the  barium  and  aluminium  salts  of 
tetraboric  and  metaboric  acids  and  calculate  the  per  cent  of  boron 
in  each  salt. 


CHAPTER  XXII 
Silicon  —  Glass 

Occurrence.  —  Silicon  does  not  occur  free  in  nature,  being 
found  almost  exclusively  as  silicon  dioxide  (Si02)  or  as  sili- 
cates. These  compounds  are  so  abundant  and  widely  dis- 
tributed that  approximately  one  fourth  of  the  earth's  crust 
is  silicon.  Sand  and  the  different  varieties  of  quartz  are  silicon 
dioxide,  while  many  rocks  are  silicates. 

Preparation  and  Properties.  —  Silicon  is  no  longer  a  rare 
element.  It  is  prepared  by  heating  a  special  mixture  of 
silicon  dioxide  and  carbon  in  an  electric  furnace  at  a  temper- 
ature which  is  carefully  regulated  to  prevent  loss  of  the  silicon 
by  volatilization  or  combination  with  the  carbon.  The 
equation  for  the  chemical  change  is  - — 

SiO2      -f      2C      =      Si        +       2  CO 

Silicon  Carbon  Silicon  Carbon 

Dioxide  Monoxide 

Thus  prepared  silicon  is  a  gray-black,  lustrous,  brittle 
solid.  It  melts  at  about  1400°  C.  and  oxidizes  to  silicon 
dioxide  when  heated  to  about  this  temperature.  The  spe- 
cific gravity  is  about  2.37.  It  is  almost  as  hard  as  quartz. 
At  high  temperatures  silicon  and  oxygen  form  silicon  dioxide ; 
with  certain  elements  silicon  forms  silicides.  Silicon  and  the 
halogens  form  volatile  compounds,  e.g.  silicon  tetrafluoride 
(SiF4)  (see  pages  365,  384).  With  sodium  hydroxide  it 
forms  sodium  silicate  (Na4SiO4)  and  hydrogen. 

Silicon  Dioxide  or  Silica,  SiO2,  is  the  most  common  com- 
pound of  silicon..  It  occurs  native  in  both  crystalline  and 

378 


SILICON 


379 


amorphous  conditions,  but  when  produced  by  a  chemical 
process  in  the  laboratory  it  is  usually  a  white  amorphous 
powder.  Sand,  gravel,  sandstone,  and  quartzite  are  almost 
wholly  silica.  It  is  an  essential  ingredient  of  many  rocks, 
as  granite  and  gneiss.  Quartz  is  silicon  dioxide.  Pure  crys- 
talline quartz  is  colorless  and  transparent,  and  is  frequently 
found  as  crystals  which  consist  usually  of  a  six-sided  prism 
with  a  six-sided  pyramid  at  one  or  both  ends ;  but  the  crystals 
are  often  distorted  or  complex  (Fig.  60).  There  are  many 
varieties  of  quartz,  which  differ 
in  color  and  structure,  due  to 
minute  impurities  or  to  the  mode 
of  formation.  Among  the  crys- 
talline varieties  are  the  clear, 
colorless  rock  crystal,  the  purple 
amethyst,  and  the  rose,  yellow, 
glassy,  milky,  and  smoky  forms. 
Varieties  imperfectly  crystalline 


FIG.  60.  —  Quartz  crystals. 


or  amorphous  are  the  waxlike  chalcedony,  the  various  forms 
of  agate  having  different  colored  layers,  the  reddish  brown 
carnelian,  the  black  and  white  onyx,  the  red  or  brown  jasper, 
the  dull  brown  or  black  flint,  and  the  brittle  chert.  Opal  is 
hydrated  silica  (SiO2 .  n  H2O) .  Petrified  or  silicified  wood 
is  largely  some  variety  of  quartz  which  has  replaced  the 
woody  fiber.  Infusorial  or  diatomaceous  earth  is  a  variety 
of  silica  consisting  of  the  shells  of  minute  organisms  called 
diatoms. 

Quartz  crystals  and  most  crystalline  varieties  of  silica 
are  hard  enough  to  scratch  glass.  They  are  insoluble  in 
water  and  acids,  except  hydrofluoric  acid,  but  are  transformed 
into  a  soluble  alkaline  silicate  when  heated  in  the  hydrox- 
ides or  carbonates  of  sodium  and  potassium.  Thus,  when 
fine  sand  is  fused  with  sodium  carbonate  the  equation  for 
the  reaction  is  — 


380  INORGANIC   CHEMISTRY 

Na2CO3       4-        SiO2        =        Na2SiO3       +       C08 

Sodium  Silicon  Sodium  Carbon 

Carbonate  Dioxide  Silicate  Dioxide 

Silica  itself  is  infusible,  except  in  the  oxyhydrogen  flame  and 
electric  furnace.  If  pure  silica  is  fused  with  certain  pre- 
cautions, the  molten  mass  can  be  shaped  into  elastic  threads, 
which  are  used  to  suspend  delicate  parts  of  electrical  instru- 
ments, and  into  tubes,  flasks,  crucibles,  etc.,  which  do  not 
crack  by  sudden  heating  and  cooling.  The  specific  gravity 
of  quartz  is  about  2.65. 

Sandstone  and  quartzite  are  used  as  building  stones,  and 
hard  sandstone  is  made  into  grindstones  and  whetstones. 
Sand  is  used  in  making  sandpaper,  glass,  porcelain,  and 
mortar.  Glass  is  roughened  and  cut  by  blowing  or  "  blast- 
ing "  fine  sand  against  it.  Many  of  the  varieties  of  quartz 
are  cut  and  polished  into  ornaments  and  gems,  e.g.  amethyst, 
opal,  and  agate.  Rock  crystal  is  used  as  the  "  diamond  " 
in  cheap  jewelry,  and  is  sometimes  cut  into  lenses  for  eye- 
glasses and  optical  instruments.  Petrified  wood  is  cut  and 
polished  into  table  tops,  mantelpieces,  and  fireplaces.  In- 
fusorial earth  is  used  to  polish  silver  ("  electro-silicon " 
being  the  commercial  name  of  one  kind)  and  in  making 
cement,  "  soluble  glass,"  dynamite,  and  refractory  brick. 

Silica  and  Plants.  —  The  ash  of  many  plants  contains  silica, 
showing  that  some  compound  of  silicon  is  assimilated  by 
the  plant  from  the  soil  —  probably  silicic  acid  or  a  soluble 
silicate  (see  below).  The  ash  of  rye  and  wheat  straws  and 
of  potato  stems  contains  from  40  to  70  per  cent  of  silica. 
Plants  like  horsetail  and  bamboo  are  rich  in  silica.  The 
silica  is  probably  not  a  plant  food  in  the  strict  sense,  but 
gives  firmness  to  the  tall  stalks,  especially  to  their  joints. 

Silicic  Acids  and  Silicates.  —  Silicon  being  a  non-metal 
forms  acids,  many  of  which  are  complex  and  known  only 


SILICON  381 

through  the  corresponding  salts.  The  two  which  are  simple 
and  best  known  are  metasilicic  acid  (H2SiO3)  and  orthosilicic 
acid  (H4SiO4).  As  stated  in  a  preceding  paragraph,  sodium 
silicate  (Na-jSiOg)  is  formed  when  silicon  dioxide  is  fused 
with  sodium  carbonate.  Now  sodium  silicate  dissolves  in 
water,  and  when  hydrochloric  acid  is  added  to  a  concentrated 
solution,  a  silicic  acid  is  precipitated  as  a  white  gelatinous 
mass.  The  precipitate  is  probably  orthosilicic  acid,  but 
when  this  precipitate  is  dried  only  metasilicic  acid  is  found 
in  the  residue,  thus :  — 

H4SiO4    =  H2SiO3    +  H2O 

Orthosilicic        Metasilicic        Water 
Acid  Acid 

The  metasilicic  acid  on  further  heating  decomposes  into 
silicon  dioxide  and  water,  thus :  — 

H2SiO3  =    SiO2    +   H2O 

Metasilicic         Silicon  Water 

Acid  Dioxide 

It  appears  then  that  these  two  silicic  acids  are  closely  related. 
Indeed  orthosilicic  acid  may  be  regarded  as  a  sort  of  parent 
of  the  other  silicic  acids,  which,  although  they  have  not  been 
isolated,  may  be  conveniently  thought  of  as  molecules  of 
orthosilicic  acid  minus  one  or  more  molecules  of  water. 
Following  out  this  relationship  we  have  for  example  the  fol- 
lowing hypothetical  silicic  acids  :  — 

2  H4SiO4  -     H2O  =  H6Si2O7,  or  Disilicic  Acid 

3  H4Si04  -4  H2O  =  H4Si3O8,  or  Trisilicic  Acid 

Colloidal  Silicic  Acid.  —  Sodium  silicate  and  hydrochloric 
acid  do  not  always  interact  as  described  above.  If  the 
sodium  silicate  solution  is  dilute,  or  the  hydrochloric  acid 
concentrated,  or  in  excess,  then  the  silicic  acid  which  is. 


382  INORGANIC  CHEMISTRY 

formed  remains  in  solution  as  colloidal  silicic  acid.  It  can- 
not be  filtered  out,  though  it  can  be  separated  b,y  dialysis 
from  the  sodium  chloride  in  the  solution.  Thus,  if  the  col- 
loidal solution  of  silicic  acid  is  placed  in  a  vessel  having  a 
bottom  of  parchment  and  hanging  in  a  larger  receptacle 
filled  with  water,  the  silicic  acid  will  be  retained  in  the 
smaller  vessel  but  the  sodium  chloride  will  pass  through  the 
parchment  into  the  water.  This  process  was  devised  by 
Graham,  who  did  the  first  work  on  colloids. 

Colloidal  Solutions.  —  In  a  colloidal  solution  the  substance  is 
suspended  as  exceedingly  fine  particles,  which  will  pass  through 
filter  paper  but  not  through  parchment  or  other  animal  membranes  ; 
neither  will  the  particles  settle,  though  they  can  be  precipitated  (or 
coagulated)  under  special  conditions.  Substances  which  form  col- 
lodial  solutions  are  called  colloids,  and  while  in  solution  are  in  the 
colloidal  condition.  Colloidal  solutions  are  not  solutions  in  the 
usual  sense,  for  unlike  true  solutions  they  show  very  little,  if  any, 
elevation  of  the  boiling  point  or  depression  of  the  freezing  point ; 
moreover,  when  a  converging  beam  of  strong  light  is  passed  through 
a  colloidal  solution,  the  path  is  brightened  by  the  particles,  whereas 
a  true  solution  remains  dark.  Sometimes  colloidal  solutions  are 
called  colloidal  suspensions. 

The  term  colloid  (from  the  Greek  word  for  glue),  which  was 
first  applied  to  sticky  substances,  now  includes  two  general  classes 
of  substances:  (1)  Those,  like  agar-agar  and  gelatin,  that  form 
jellylike  masses  on  cooling  or  concentration  ;  and  (2)  those,  like  gold 
and  arsenious  sulphide,  that  coagulate  (i.e.  precipitate  upon  the 
addition  of  an  electrolyte).  Certain  colloids  are  precipitated  merely 
by  heating.  Many  colloids  carry  electrical  charges.  Arsenious 
sulphide  (As2S3),  silver  chloride  (AgCl),  and  certain  metals  (Ag, 
Cu,  Pt)  are  negative,  while  ferric  hydroxide  (Fe(OH)3)  and  many 
basic  substances  are  positive;  starch  and  gelatin  are  neutral. 
Charged  colloids  are  precipitated  by  oppositely  charged  ions  and 
colloids.  Thus,  colloidal  arsenious  sulphide  is  precipitated  by  hy- 
drogen ion  (e.g.  from  hydrochloric  acid)  and  also  by  colloidal  ferric 
hydroxide ;  so  also  negative  metaphosphoric  acid  and  positive 
(usually)  albumin  precipitate  one  another  (see  page  399). 

Coagulation  of  colloids  can  be  retarded  or  prevented  by  adding 


SILICON  383 

a  protective  colloid.  Thus,  gelatin  is  often  used  to  render  col- 
loidal silver  bromide  more  stable,  e.g.  in  photographic  plates.  It 
is  supposed  that  the  protective  colloid  forms  a  film  around  the 
other  colloid  and  thereby  prevents  diffusion. 

Silicates  are  the  salts  of  silicic  acids,  though  their  cor- 
responding acids  have  not  been  isolated  in  most  cases.  So- 
dium and  potassium  silicates  are  salts  of  the  well-known 
metasilicic  acid  (H2SiO3) .  They  are  the  only  silicates  soluble 
in  pure  water,  and  the  thick,  sirupy  solution  of  each  (or 
both)  is  called  water  glass.  It  finds  extensive  use  in  the 
manufacture  of  soap,  certain  cements,  artificial  stone,  and 
fireproof  materials.  As  already  stated  many  rocks  and  min- 
erals are  silicates  and  make  up  a  large  part  of  the  earth's 
crust.  The  following  list  shows  the  relations  of  a  few 
silicates  to  their  acids :  — 

f  Wollastonite,  CaSiO8 
Metasilicates  I  ^        ...     -.,  r?.^ 

fo.  u      t  TTQTk\         Enstatite,  Mgbi03 

(SaltS  Of   H2SlO3)  -r,          n      -r,       Al/crn\ 

[  Beryl,  Be3Al2(SiO3)6 

f  Zircon.  ZrSiO4 
Orthosmcates  I  ,r    ,.      „  A1  /0.~s     „  ~ 

rc«it«  nf  TT  <3in  ^     I  Kaolm;  H2AJ2(SiO4)2 .  H2O 
(baits  01  ±140104)         ~,.  .       ,,    0..~     -r,   p..^ 
[  Ohvme,  Mg2SiO4 .  Fe2SiO4 

Disilicate  Serpentine,  Mg3Si2O7 .  2  H2O 

(Salt  of  H6Si2O7) 

Trisilicate  Orthoclase,  KAlSi3O8 

(Salt  of  H4Si308) 

Other  silicates  are  mica,  hornblende,  augite,  slate,  talc,  lava, 
feldspars  related  to  orthoclase,  asbestos,  garnet,  and  tour- 
maline. The  most  abundant  are  the  silicates  of  calcium, 
aluminium,  magnesium,  potassium,  sodium,  and  iron. 

Silicic  acid  is  a  feeble  acid.  It  does  not  redden  blue  litmus 
nor  liberate  hydrogen  when  added  to  magnesium.  Never- 
theless it  is  properly  called  an  acid  because  it  forms  salts. 
These  salts,  which  we  have  already  seen  make  up  a  large  part 


384  INORGANIC  CHEMISTRY 

of  the  earth's  crust,  are  slowly  decomposed  by  carbonic  acid, 
i.e.  by  the  joint  action  of  the  carbon  dioxide  and  water  vapor 
in  the  atmosphere.  This  disintegration  of  the  silicates  is 
called  weathering.  , 

The  water  of  many  hot  springs,  as  in  the  Yellowstone 
Park,  contains  alkaline  silicates ;  and  when  the  solution  comes 
to  the  surface,  some  of  the  silicate  is  decomposed  by  the 
carbon  dioxide  in  the  air,  and  the  silica  is  deposited  around 
the  spring  in  beautiful  forms  called  geyserite  or  siliceous 
sinter. 

Silicon  Tetrafluoride,  SiF4,  is  a  colorless  gas  which  has  a 
pungent,  suffocating  odor.  It  is  formed  when  hydrofluoric 
acid  interacts  with  silicon  dioxide  or  silicates.  Thus,  with 
silicon  dioxide  the  equation  is  — 

SiO2     +      4HF       =       SiF4     +     2H2O 

Silicon          Hydrofluoric  Silicon 

Dioxide  Acid  Tetrafluoride 

Silicon  tetrafluoride  forms  fumes  in  moist  air  and  interacts 
readily  with  water,  thus :  — 

3SiF4   +     4H2O     =     H4SiO4     +     2  H2SiF6 

Silicon  Silicic  Hydrofluosilicic 

Tetrafluoride  Acid  Acid 

The  hydrofluosilicic  acid  (sometimes  called  simply  fluosilicic 
acid)  remains  in  solution,  while  the  silicic  acid  is  precipitated. 
The  formation  of  the  white  gelatinous  silicic  acid  when  the 
gases  from  the  interaction  of  hydrofluoric  acid  and  a  com- 
pound of  silicon  are  led  into  water  is  often  used  as  a  test  for 
silicon. 

Siloxicon  is  the  commercial  name  of  a  highly  refractory 
substance  produced  by  heating  a  mixture  of  silicon  dioxide 
and  carbon  to  about  2500°  C.  in  a  special  form  of  an  electric 


SILICON  385 

furnace.  It  is  not  a  definite  compound,  but  varies  in  com- 
position from  Si2C2O  to  Si7C7O.  It  is  a  gray-green,  granular 
powder  which  can  be  readily  shaped  into  bricks,  linings,  and 
other  forms  of  refractory  articles. 

Silicides  are  compounds  of  silicon  and  other  elements. 
Carborundum  or  carbon  silicide  is  the  most  important. 

Carborundum,  SiC,  is  a  crystalline  compound  consisting 
solely  of  silicon  and  carbon.  It  varies  in  color  from  white 
to  emerald  green  and  is  sometimes  iridescent.  It  is  ex- 
tremely hard,  being  nearly  as  hard  as  diamond.  The  specific 
gravity  is  about  3.  Acids  do  not  affect  it,  but  it  is  decom- 
posed by  fusing  it  with  potassium  hydroxide  and  other  alkalies. 

Its  extreme  hardness  has  led  to  its  application  as  an  abra- 
sive, and  large  quantities  are  made  into  a  great  variety  of 
grinding  wheels,  whetstones,  and  polishing  cloths. 

Carborundum  is  manufactured  by  fusing  a  mixture  of 
sand  and  coke  in  an  electric  furnace  constructed  on  the 
resistance  type  (Fig.  61).  It  is  essentially  an  oblong  box 
with  permanent  ends  and  loosely  built  sides.  Each  end  is 
provided  with  a  heavy  metal  plate.  The  wires  for  the 
electric  current  are  attached  to  the  outer  ends  of  these 
plates,  while  the  huge  carbon  electrodes  fit  into  the  inner 
ends,  and  project  into  the  furnace.  A  cylindrical  mass  of 
granulated  coke  makes  an  electrical  connection  between  the 
electrodes.  The  mixture  of  sand  and  coke  (to  which  salt  and 
sawdust  are  added  to  contribute  to  the  fusion  and  porosity) 
is  packed  around  this  core  inside  the  box.  The  heat  gen- 
erated by  the  resistance  of  the  carbon  core  to  the  passage  of 
the  powerful  current  of  electricity  produces  a  chemical 
change  essentially  as  follows :  — 

SiO2   +    3C      =      SiC     -f     2  CO 

Sand          Carbon        Carborundum        Carbon 

Monoxide 


386 


INORGANIC   CHEMISTRY 


The  change  is  due  solely  to  the  intense  heat,  i.e.  it  is  an 
electrothermal,  not  an  electrolytic  change.  When  the  opera- 
tion is  over  and  the  furnace  is  cool,  the  side  walls  are  pulled 
down,  and  the  carborundum  is  removed.  The  purest  grade 


GLASS  387 

is  found  around  the  core.  The  product  is  crushed,  treated 
with  sulphuric  acid  to  remove  the  impurities,  washed,  dried, 
and  graded  according  to  the  size  of  the  particles. 

Miscellaneous.  —  Silicon  is  a  non-metallic  element  and 
belongs  to  the  carbon  family  in  the  fourth  (IV)  periodic 
group.  Certain  physical  properties  suggest  a  metallic  char- 
acter, but  chemically  silicon  is  very  closely  related  to  carbon. 
Both  have  allotropic  modifications  and  form  analogous  com- 
pounds, e.g.  CO2  and  SiO2,  CH4  and  SiH4.  Both  form  many 
compounds  of  great  importance,  so  that  we  might  conven- 
iently regard  silicon  as  the  chief  element  in  the  mineral  king- 
dom, just  as  carbon  is  in  the  organic  realm. 

Silicon  has  the  valence  of  four  in  its  compounds. 

GLASS 

Glass  is  an  amorphous,  more  or  less  transparent  solid.  It 
is  a  homogeneous  mixture  of  silicates  with  an  excess  of 
silica.  Glass  is  not  made  by  mixing  silicates,  but  by  fusing 
a  mixture  of  sand,  an  alkali,  and  a  calcium  or  lead  compound. 
The  alkali  is  potassium  carbonate  (K2C03)  or  sodium  car- 
bonate (Na2CO3),  though  sodium  sulphate  and  sodium  nitrate 
are  used  in  some  cases  as  auxiliary  substances;  the  calcium 
compound  is  limestone  (CaCO3)  or  lime  (CaO) ;  and  the  lead 
compound  is  litharge  (PbO)  or  red  lead  (Pb3O4).  Besides 
these  fundamental  ingredients,  small  quantities  of  other 
substances  are  used,  e.g.  (I)  broken  glass  (called  cullet), 
which  lowers  the  melting  point  of  the  mixture;  (2)  arsenic 
trioxide  (As2O3),  which  destroys  carbonaceous  impurities; 
(3)  carbon,  which  lowers  the  melting  point  when  sodium 
sulphate  is  used  and  likewise  imparts  a  color  from  straw  to 
amber ;  and  (4)  manganese  dioxide  (Mn02) ,  which  neutralizes 
the  green  color  caused  by  iron  compounds  (often  present  in 
impure  materials). 


388  INORGANIC   CHEMISTRY 

The  process  consists  in  melting  a  carefully  prepared  mix- 
ture of  the  proper  ingredients  in  a  refractory  fire-clay  pot. 
The  heat  is  often  obtained  by  burning  gas  —  manufactured 
or  natural.  During  the  melting,  gases  escape,  and  the  im- 
purities, which  rise  to  the  surface  as  a  scum,  are  removed. 
The  molten  mass  is  allowed  to  cool  until  it  has  the  proper 
consistency.  A  portion  is  then  collected  as  a  soft  ball  on 
the  end  of  an  iron  tube  and  brought  to  the  desired  shape, 
either  by  forcing  it  into  a  mold  or  by  blowing  into  the  tube 
and  simultaneously  manipulating  the  plastic  mass  by  twist- 
ing and  swinging.  The  details  of  the  procedure,  however, 
vary  with  the  article  being  made.  Many  objects,  such  as 
tumblers  and  small  dishes,  are  now  made  by  pressing  the 
plastic  glass  with  a  die  or  by  blowing  it  into  a  mold. 
Fruit  jars,  bottles,  and  lamp  chimneys  are  blown  by 
machinery. 

All  glass  must  be  cooled  slowly  to  prevent  brittleness. 
This  operation  is  called  annealing,  and  is  accomplished  by 
passing  the  objects  slowly  through  a  furnace  in  which  the 
temperature  is  gradually  lowered. 

There  are  four  possible  kinds  of  glass  and  many  varieties 
of  each.  Their  properties  depend  upon  the  proportion  of  the 
ingredients,  and  each  kind  may  be  made  to  approach  the 
others  in  properties  by  varying  these  proportions.  Arranged 
in  the  order  of  their  fusibility  and  beginning  with  the  soft- 
est, the  four  kinds  of  glass  are:  (1)  Sodium-lead  glass,  (2) 
potassium-lead  glass,  (3)  sodium-calcium  glass,  and  (4) 
'potassium-calcium  glass.  Flint  glass  is  a  lead  glass;  it  is 
lustrous,  refracts  light  to  a  high  degree,  and  is  made  into 
ornaments,  lenses  for  optical  instruments,  and  also  into 
shades  for  electric  and  gas  lights.  Cut  glass  objects  are 
made  from  flint  glass  by  means  of  simple  grinding  and  pol- 
ishing machinery.  Window,  plate,  crown,  table,  and  bottle 
glass  is  a  sodium-calcium  glass ;  it  is  sometimes  called 


GLASS  389 

soda  glass  or  soft  glass  to  distinguish  it  from  the  potassium- 
calcium  glass,  which  is  hard. 

Window  glass  is  made  by  blowing  a  lump  of  glass  into  a 
hollow  globe  and  then  into  a  cylinder;  this,  on  being  opened 
at  both  ends  and  cut  lengthwise,  spreads  out  flat.  Plate 
glass  is  made  by  pouring  the  molten  glass  upon  a  large  table, 
rolling  it  with  a  hot  iron  roller,  and  subsequently  grinding  and 
polishing  it  until  the  surfaces  are  parallel.  Plate  glass  is 
used  for  large  windows  and  for  mirrors,  but  considerable 
rough  plate  is  used  for  skylights  and  floors.  Crown  glass  is 
a  superior  quality  of  window  glass.  It  has  a  brilliant  sur- 
face and  is  used  as  "bull's-eyes"  in  decorative  windows  and 
as  lenses  for  optical  instruments  (in  conjunction  with  flint 
glass).  Most  chemical  glassware  is  sodium-calcium  glass. 
It  therefore  softens  when  heated  and  the  flame  becomes 
yellow  from  the  sodium.  Bohemian  or  hard  glass  is  a  potas- 
sium-calcium glass.  It  is  much  harder  than  the  other  kinds 
and  is  used  in  making  chemical  apparatus  designed  to  with- 
stand great  heat.  Soft  glass  is  slightly  soluble  in  water,  but 
hard  glass  is  less  so,  hence  special  varieties  of  hard  glass  are 
often  made  into  apparatus  which  resists  the  solvent  action 
of  water  and  chemical  reagents;  Jena  glass  is  one  variety. 

Colored  glass  is  made  by  adding  different  substances  to 
the  mixture.  Iron,  chromium,  and  certain  copper  com- 
pounds make  it  green,  the  green  color  of  many  bottles  and 
fruit  jars  being  due  to  iron  compounds  in  the  impure  ma- 
terials used ;  cobalt  compounds  produce  different  shades  of 
blue;  manganese  dioxide  gives  a  pink  or  a  violet  color; 
yellow  is  produced  by  charcoal,  sulphur,  uranium  compounds, 
or  silver ;  the  deep  red  glass  so  extensively  used  in  lanterns 
is  usually  colored  by  selenium  compounds ;  milky  glass  is 
made  by  adding  calcium  phosphate,  fluor  spar,  or  cryolite ; 
stained  glass  is  ordinary  glass  to  which  fusible  pigments  are 
applied  with  a  brush  and  then  fixed  by  heat ;  iridescent  glass 


390  INORGANIC  CHEMISTRY 

is  made  by  secret  processes,  though  it  is  known  that  one 
consists  in  exposing  a  special'  variety  of  absorbent  glass  to 
the  vapors  of  metallic  oxides. 

PROBLEMS  AND  EXERCISES 

1.  How  can  SiO2  be  transformed  into  H2Si03?     How  many 
grams  of  SiO2  are  needed  for  75  gm.  of  H2SiO3? 

2.  How  can  SiO2  be  transformed  into  H2SiF6?     How  much  SiO2 
is  needed  for  100  gm.  of  H2SiF«? 

3.  How  much  hydrofluosilicic  acid  can  be  made  by  the  inter- 
action of  water  and  a  metric  ton  of  silicon  tetrafluoride  ? 

4.  Calculate  the  atomic  weight  of  silicon  from  the  following 
data :    (a)  2.621  gm.  of  silicon  tetrachloride  (SiCl4)  required  6.6445 
gm.  of  silver  for  precipitation  of  the  chlorine ;    the  atomic  weight 
of  silver  was  accepted  as  107.88.     (6)  95.52367  gm.  of  silicon  tetra- 
bromide  (SiBr4)  yielded  16.56868  gm.  of  silicon  dioxide  ;  the  atomic 
weights  of  oxygen  and  bromine  were  accepted  as  16  and  79.92 
respectively. 

5.  Write  the  formulas  of  the  following  compounds  by  applying 
the  principle  of  valence  (see  Chapter  XIV) :  Silicon  iodide,  hydrogen 
silicide,  silicon  sulphide,  carbon  silicide. 

6.  Calculate  the  per  cent  of  silicon  in  (a)  calamine,  Zn2SiO4, 
(6)  chrysocolla,  CuSi03,  (c)  analcite,  Na2Al2Si4Oi2,  (d)  heulandite, 
CaAl2Si6Oifl,  (e)  beryl,  Be3Al2Si«Oi8  (Be  =  9),  (/)  phenacite,  Be2SiO4, 
(0)    garnet,    Ca3Al2Si3Oi2,     (h)    muscovite,    KAlSiO4,     (i)    olivine, 
Mg2SiO4 .  FeS04,  (7)  zircon,  ZrSiO4  (Zr  =  90.6). 

7.  How  much  silicon  can  be  made  (a)  by  reducing  a  metric  ton 
of  sand  (90  per  cent  pure)  with  C?   (6)  from  119  gm.  of  potassium 
silico-fluoride?     (Equation  is  K2SiF6  +  4  K  =  Si  +  6  KF.) 

8.  How  much  metasilicic  acid  (H2Si03)  can  be  made  by  the  in- 
teraction of  water  and  a  metric  ton  of  silicon  tetrafluoride  ? 

9.  Suppose  opal  is  SiO2 .  10  H2O ;  calculate  its  per  cent  of  (a) 
SiO2,  (6)  Si,  (c)  H2O. 

10.  (a)  How  much  metasilicic  acid  can  be  made  from  a  metric 
ton  of  orthosilicic  acid  ?     (6)  How  much  SiO2  from  a  metric  ton  of 
metasilicic  acid? 

11.  Calculate  the  simplest  formulas  corresponding  to  (a)  Si  = 
35.897,  H  =  2.564,  O  =  61.538;  (6)  Si  =  29.166,  H  =  4.166,  O  =  66.666. 
What  is  the  name  of  each  compound? 


CHAPTER   XXIII 
Phosphorus,  Arsenic,  Antimony,  and  Bismuth 

PHOSPHORUS,  arsenic,  antimony,  and  bismuth,  together 
with  nitrogen,  belong  to  the  nitrogen  family  in  the  fifth  (V) 
periodic  group  of  elements. 

PHOSPHORUS 

Occurrence.  —  Free  phosphorus  is  not  found  in  nature. 
But  phosphates  are  numerous  and  abundant,  the  most 
common  being  phosphorite  ("  phosphate  rock,"  impure 
Ca3(P04)2)  and  apatite  (Ca5F(PO4)3).  Approximately  .1 
per  cent  of  the  earth's  crust  is  phosphorus.  Phosphates  are 
present  in  fertile  soils  and  some  iron  ores.  Phosphorus 
compounds  are  essential  constituents  of  seeds,  and  also  of 
the  brain,  nerves,  muscles,  and  bones  of  animals. 

Phosphorus  was  discovered  in  1669  by  Brand,  a  German  alchemist, 
who  obtained  it  by  heating  a  certain  kind  of  animal  matter.  Scfyeele, 
in  1771,  extracted  it  from  bones. 

Preparation.  —  Phosphorus  is  prepared  industrially  by 
two  processes:  (1)  In  the  older  process  bone  ash  (which  is 
over  80  per  cent  calcium  phosphate)  or  a  native  phosphate  is 
finely  ground  and  mixed  in  large  vats  with  enough  sulphuric 
acid  to  produce  the  following  change :  — 

Ca3(P04)2    +    3H2SO4       =      2H3P04        +       3  CaSO4 

Calcium  Sulphuric  Phosphoric  Acid  Calcium 

Phosphate  Acid  (Ortho-)  Sulphate 

The  insoluble  calcium  sulphate  is  removed  by  filtering  the 
mixture  through  cinders.     The  phosphoric  acid  solution  is 

391 


392 


INORGANIC  CHEMISTRY 


concentrated,  mixed  with  sawdust,  coke,  or  charcoal,  and 
dried,  being  changed  thereby  into  metaphosphoric  acid  ac- 
cording to  the  equation :  — 

H3PO4  =    ,       HP03        -I-        H20 

Phosphoric  Acid  Phosphoric  Acid 

(Ortho-)  (Meta-) 

The  dried  mass  is  heated  to  a  high  temperature  in  clay 
retorts  arranged  in  tiers  (Fig.  62),  the  change  thus  produced 
being  substantially  — 

4HPO3    +     12  C       =       P4       -h       2H2      -f      12  CO 

Phosphoric  Acid        Carbon  Phosphorus        Hydrogen  Carbon 

(Meta-)  Monoxide 

The  phosphorus  distills  as  a  vapor  through  a  pipe  into  a 

trough  of  water,  where  it  collects 


FIG.  62.  — Apparatus  for  the  manu- 
facture of  phosphorus  by  the  old 
method  (final  stage). 


FIG.  63.  —  Electric  furnace  for 
the  manufacture  of  phos- 
phorus. The  raw  materials  in- 
troduced at  A  are  fed  in  by 
the  screw  B,  the  phosphorus 
vapor  escapes  at  C,  and  the 
slag  is  drawn  off  at  D.  The 
electrodes  are  E,  E. 


as  a  heavy  liquid.     (2)  By  a   new  process   phosphorus  is 
manufactured  in  an  electric  furnace  (Fig.  63).     A  mixture 


PHOSPHORUS  393 

of  phosphate,  carbon,  and  sand  is  fed  continuously  into  a 
furnace  provided  with  an  outlet  pipe  near  the  upper  part 
through  which  the  phosphorus  vapor  passes  into  a  condenser. 
The  residue  is  drawn  off  as  a  slag  at  the  bottom  of  the  fur- 
nace. The  process  is  an  electrothermal  one,  the  essential 
equation  for  the  chemical  change  being  — 

2  Ca3(PO4)2  +  6  SiO2  +  100     =     P4    +    10  CO  +  6  CaSi03 

Calcium  Sand         Carbon      Phosphorus       Carbon  Calcium 

Phosphate  Monoxide         Silicate 

Each  method  gives  a  black  product,  which  is  purified  by 
redistillation  in  an  iron  retort,  or  by  oxidation  under  water 
with  sulphuric  acid  and  potassium  dichromate;  finally  it  is 
pressed  through  canvas  bags  and  molded  into  sticks. 

Properties.  —  Phosphorus  has  two  allotropic  modifica- 
tions, —  yellow  (waxy)  or  ordinary  and  red.  Ordinary 
phosphorus  when  freshly  prepared  is  a  yellow,  translucent 
solid,  but  the  color  deepens  by  exposure  to  light.  At  ordi- 
nary temperatures  it  is  like  wax,  but  at  low  temperatures 
it  is  brittle.  Under  water  it  melts  at  44°  C.  Exposed  to 
the  air,  it  immediately  gives  off  white  fumes,  and  at  34°  C. 
takes  fire  and  burns  with  a  brilliant  flame,  the  main  product 
being  phosphorus  pentoxide  (P205).  It  is  luminous  in 
moist  air,  as  may  be  easily  seen  by  rubbing  a  phosphorus- 
tipped  match  in  a  dark  room.  This  property  gave  the  ele- 
ment its  name  (from  the  Greek  word  phosphoros,  light 
bringer).  The  ease  with  which  it  ignites  makes  phosphorus 
dangerous  to  handle.  Phosphorus  is  kept  beneath  water, 
and  should  never  be  handled  or  cut  unless  so  covered. 
Burns  from  it  are  severe  and  hard  to  heal.  It  is  very  poison- 
ous, and  the  workmen  in  factories  where  phosphorus  is  used 
are  liable  to  contract  a  dreadful  disease,  which  rots  the  bones. 
Phosphorus  is  nearly  insoluble  in  water,  but  dissolves  in 
carbon  disulphide  and  slightly  in  sodium  hydroxide  solution. 


394  INORGANIC  CHEMISTRY 

Yellow  phosphorus  has  a  faint  but  characteristic  odor, 
which  may  be  easily  detected  by  smelling  of  a  phosphorus- 
tipped  match.  Red  phosphorus  is  made  by  heating  ordi- 
nary phosphorus  to  about  250°  C.  in  a  closed  vessel  freed 
from  air.  Conversely,  if  red  phosphorus  is  distilled  and  the 
vapor  condensed  quickly,  the  yellow  variety  is  obtained. 
This  red  modification  of  phosphorus  is  a  dark  red  powder, 
though  sometimes  it  is  a  brittle  mass.  It  is  opaque  and 
odorless,  does  not  glow  or  take  fire  when  exposed  to  the 
air,  and  does  not  ignite  until  heated  to  260°  C.  It  is  not 
poisonous,  and  does  not  dissolve  in  carbon  disulphide.  It 
can  be  handled  without  danger  and  need  not  be  kept  be- 
neath water.  Obviously  it  is  the  less  active  variety  of  the 
element.  The  specific  gravity  of  red  phosphorus  is  about 
2.2  and  that  of  the  yellow  form  is  about  1.83. 

Certain  rat  and  bug  poisons  contain  yellow  phosphorus, 
but  most  of  the  phosphorus  of  commerce  is  consumed  in  the 
manufacture  of  matches  (see  below). 

The  vapor  density  of  both  yellow  and  red  phosphorus  up 
to  approximately  1500°  C.  corresponds  to  the  molecular 
weight  128.  Since  the  atomic  weight  is  31,  the  molecular 
formula  is  P4.  At  higher  temperatures  partial  dissociation 
occurs.  The  formula  of  dissolved  phosphorus  is  also  P4. 

Oxides  of  Phosphorus.  —  The  two  important  oxides  are 
phosphorous  or  the  trioxide  (P2O3  or  P4O6)  and  phosphoric 
or  the  pentoxide  (P205  or  P4O10).  Phosphorous  oxide  is  a 
white  solid  formed  by  the  slow  oxidation  of  phosphorus  or 
by  burning  phosphorus  in  a  limited  supply  of  air.  It  has 
the  odor  of  phosphorus  and  is  poisonous.  Warmed  in  the 
air,  it  changes  into  the  pentoxide.  It  unites  with  water  to 
form  phosphorous  acid,  thus :  — 

P2O3         -f-         3H2O       =       2H8PO8 

Phosphorous  Phosphorous 

Oxide  Acid 


PHOSPHORUS  395 

Phosphoric  oxide  is  a  white,  snowlike  solid  formed  by  burn- 
ing phosphorus  in  an  abundant  supply  of  air.  It  is  very  de- 
liquescent and  is  often  used  in  the  laboratory  to  dry  gases. 
It  combines  vigorously  with  cold  water,  forming  metaphos- 
phoric  acid,  thus :  — 


P205        +        H20       =         2HPOS 

aosphoric  Metaphosphc 

Oxide  Acid 


Acids  and  Salts  of  Phosphorus.  —  There  are  three  phos- 
phoric acids,  —  orthophosphoric  (H3PO4),  metaphosphoric 
(HPO3),  and  pyrophosphoric  (H4P2O7).  Phosphorous  acid 
(H3PO3)  and  hypophosphorous  (H3PO2)  are  less  important 
compounds. 

Orthophosphoric  Acid  is  a  by-product  in  the  manufacture  of 
phosphorus  from  bone  ash  (see  above);  it  can  be  prepared 
by  oxidizing  red  phosphorus  with  dilute  nitric  acid,  or  by 
dissolving  phosphorus  pentoxide  in  hot  water,  thus :  — 

P2O5       +       3H2O      =       2H3P04 

Phosphorus  Orthophosphoric 

Pentoxide  Acid 

It  is  a  white  deliquescent  solid  which  is  very  soluble  in  water. 

Metaphosphoric  Acid  is  formed  by  heating  orthophosphoric 
acid  to  a  high  temperature,  thus :  — 

H3P04  HPO3       +       H2O 

Orthophosphoric  Metaphosphoric 

Acid  Acid 

It  can  be  formed  by  dissolving  the  pentoxide  in  cold  water, 
thus : — 

P2O5  -f  H2O  =  2  HPO3 

At  ordinary  temperature  it  is  a  glassy  solid,  and  is  called 
glacial  phosphoric  acid,     It  dissolves,  readily  in  water,  and 


396  INORGANIC  CHEMISTRY 

the  solution  changes  into  orthophosphoric  acid  —  slowly  in 
the  cold,  rapidly  when  boiled. 

Pyrophosphoric  Acid  is  formed  by  heating  orthophosphoric 
acid  or  one  of  its  salts  to  200°-300°  C.,  thus:  — 

2H3P04     =       H4P207     +     H20 

Orthophosphoric      Pyrophosphoric 
Acid  Acid 

A  sodium  salt  (HNa2P04)  of  the  ortho-acid  is  usually  used. 
This  acid  is  an  amorphous,  glassy  (but  sometimes  crystalline) 
.solid,  readily  soluble  in  water. 

The  acids  of  phosphorus  just  discussed  form  salts  called 
phosphates,  many  of  which  are  found  as  minerals,  especially 
phosphates  of  calcium,  aluminium,  and  magnesium.  The 
bones  of  animals  and  ashes  of  plants  contain  calcium  and 
magnesium  phosphates.  Orthophosphoric  acid  is  tribasic, 
and  its  salts,  which  are  usually  called  simply  phosphates, 
are  numerous.  They  are  known  as  primary,  secondary, 
and  tertiary  phosphates,  according  as  one,  two,  or  three 
atoms  of  hydrogen  are  replaced.  The  most  important  is  the 
normal  calcium  salt  Ca3(PO4)2,  which  has  already  been  de- 
scribed. Hydrogen  disodium  phosphate  (HNa2PO4)  is  the 
commercial  sodium  phosphate;  it  is  a  secondary  phosphate. 
This  salt  and  hydrogen  sodium  ammonium  phosphate,  or 
microcosmic  salt  (HNa(NH4)PO4),  are  used  in  chemical 
analysis.  The  "acid  phosphate"  sold  as  a  beverage  is  a 
solution  of  one  or  more  acid  calcium  phosphates  (HCaPO4 
and  H4Ca(P04)2).  Metaphosphates  are  formed  by  heating 
primary  (or  mono-)  sodium  phosphates,  thus :  — 

H2NaPO4      =     NaPO3     +     H2O 

Primary  Sodium         Sodium  Meta- 
Phosphate  phosphate 

Microcosmic  salt  when  fused  also  forms  a  metaphosphate 
(NaPO3)  owing  to  the  loss  of  water  and  ammonia  (NH3) .  The 


PHOSPHORUS  397 

glassy  residue  is  called  a  phosphorus  bead  and  like  the  borax 
bead  assumes  different  colors  when  heated  with  metallic 
oxides.  Pyrophosphates  (of  which  only  two  classes  are 
known)  are  formed  by  heating  secondary  (or  di-)  phosphates, 
thus : — 

2HNa2PO4    =    Na4P2O7    +  H2O 

Disodium  Sodium  Pyro- 

Phosphate  phosphate 

Hypophosphites  are  the  salts  of  hypophosphorous  acid 
and  are  produced  by  treating  phosphorus  with  an  alkali. 
They  are  often  used  as  medicines. 

Orthophosphoric  acid  dissociates  mainly  into  the  ions 
H+  and  H2POr.  The  disodium  phosphate  (HNa2PO4) 
dissociates  into  the  ions  2  Na+  and  HPO4~  ".  Its  solution 
is  slightly  alkaline  because  hydroxyl  ions  are  left  when  the 
hydrogen  ion  of  the  slightly  dissociated  water  forms  the 
H2P04-ion  with  the  HPO4-ion,  the  latter  ion  having  only  a 
very  slight,  if  any,  tendency  to  dissociate;  the  simple  ionic 
equation  for  the  hydrolysis  is  — 

H+  +  OH-  -f  HPOr  "  ->  OH-  +  H2POr 

Tests  for  Phosphoric  Acids  and  Phosphates.  —  Phosphates 
can  be  distinguished  by  silver  nitrate.  Orthophosphates 
give  yellow  silver  phosphate  (Ag3PO4),  metaphosphates  give 
white  silver  metaphosphate  (AgP03),  pyrophosphates  give 
white  silver  pyrophosphate  (Ag4P2O7) ;  all  dissolve  in  am- 
monium hydroxide.  Metaphosphoric  acid  coagulates  a 
solution  of  albumin  (e.g.  white  of  egg),  but  orthophosphoric 
and  pyrophosphoric  acids  do  not.  Orthophosphoric  acid 
and  its  salts  precipitate  yellow  ammonium  phosphomolyb- 
date  from  an  excess  of  a  solution  of  ammonium  molybdate. 

Other  Compounds  of  Phosphorus.  —  Phosphine  (PH3)  is 
analogous  to  ammonia  (NH3),  though  it  is  not  alkaline.  It 


398 


INORGANIC   CHEMISTRY 


is  made  by  heating  sodium  (or  potassium)  hydroxide  with 
phosphorus.  It  is  poisonous,  and  has  a  disagreeable  odor; 
it  usually  takes  fire  spontaneously  in  the  air  owing  to  the 
presence  of  an  inflammable  compound  of  phosphorus  and 
hydrogen.  Phosphine  itself  does  not  take  fire  spontaneously. 
It  combines  with  other  substances,  forming  phosphonium 
compounds,  which  are  analogous  to  ammonium  compounds, 
e.g. :  — 

PH3    +      HI       =      PHJ 


Phosphine 


Hydriodic 
Acid 


Phosphonium. 
Iodide 


Phosphorus  trichloride  (PC13)  is  a  disagreeable  smelling  liquid, 
made  by  the  combustion  of  dry  chlorine  and  phosphorus; 
and  phosphorus  pentachloride  (PC15)  is  a  greenish  solid  made 
by  passing  chlorine  into  a  vessel  containing  the  trichloride. 
Both  trichloride  and  pentachloride  interact  readily  with 
water,  forming  phosphorus  compounds  and  hydrochloric 
acid,  thus:  — 


PC13 

Phosphorus 
Trichloride 

PC15 

Phosphorus 
Pentachloride 

PC15 

Phosphorus 
Pentachloride 


3H2O 

Water 

H20 

Water 

4H2O 

Water 


H3PO3 

Phosphorous 
Acid 

POC13 

Phosphorus 
Oxychloride 

H3P04 

Phosphoric 
Acid 

+      3HC1 

Hydrochloric 
Acid 

+        2HC1 

Hydrochloric 
Acid 

-f-      5  HC1 

Hydrochloric 
"    Acid 

If  either  chloride  is  exposed  to  moist  air,  white  fumes  are 
formed  owing  to  the  liberation  of  hydrogen  chloride  (HC1). 

When  phosphorus  pentachloride  is  heated,  it  sublimes 
without  melting,  and  under  special  conditions  of  temper- 
ature and  pressure  it  becomes  a  vapor.  The  molecular 
weight  determined  from  the  density  of  this  vapor  at  about 
300°  C.  is  only  about  half  the  calculated  value,  i.e.  104  in- 
stead of  208.5.  Examination  of  this  vapor  shows  that  it  is 


PHOSPHORUS  399 

not  phosphorus  pentachloride,  but  almost  entirely  a  mixture 
of  phosphorus  trichloride  vapor  and  chlorine  gas.  The  color 
is  greenish  and  these  two  components  can  be  separated  by 
diffusion.  These  facts  mean  that  phosphorus  pentachloride 
when  heated  dissociates  into  phosphorus  trichloride  and 
chlorine.  The  reaction  is  reversible  and  the  equation  may 
be  written  thus :  — 

PC15  ^t  PC13  +  C12 

At  300°  C.  equilibrium  is  maintained  by  about  3  per  cent  of 
phosphorus  pentachloride  and  97  per  cent  of  the  trichloride 
and  chlorine.  This  equilibrium  may  be  displaced  and  the 
reaction  sent  in  the  reverse  direction  by  increasing  the  con- 
centration of  one  of  the  reaction  products,  i.e.  by  adding 
phosphorus  trichloride  or  chlorine  to  the  tube  in  which  the 
equilibrium  prevails.  As  a  matter  of  fact,  when  phosphorus 
pentachloride  is  vaporized  in  a  vessel  containing  an  excess  of 
the  trichloride  vapor,  the  dissociation  of  the  pentachloride  is 
reduced  to  a  minimum,  for  the  vapor  density  of  the  penta- 
chloride then  corresponds  to  a  molecular  weight  of  209,  the 
calculated  value  being  208.5.  (See  Chemical  Equilibrium, 
Chapter  XI.) 

Matches.  —  Phosphorus  until  1913  was  chiefly  used  in  the  manu- 
facture of  matches.  Now  a  prohibitive  tax  compels  the  use  of  a 
substitute,  which  is  usually  a  phosphorus  sulphide  (P4S3).  In 
making  common  matches  one  end  of  the  match  stick  is  first  dipped 
into  melted  sulphur  or  paraffin  and  then  into  the  "  phosphorus  mix- 
ture." The  latter  consists  of  different  proportions  of  phosphorus 
sulphide,  manganese  dioxide  (or  other  oxidizing  substance),  glue, 
and  a  little  coloring  matter.  By  rubbing  them  on  a  rough  surface 
enough  heat  is  generated  to  cause  the  phosphorus  to  unite  with  the 
oxygen  of  the  oxidizing  agent,  and  the  heat  thereby  produced  sets 
fire  to  the  sulphur  or  paraffin,  and  this  in  turn  kindles  the  wood. 
In  safety  matches  the  head  is  usually  a  colored  mixture  of  anti- 
mony sulphide,  potassium  chlorate,  and  glue,  while  the  rubbing 
surface  is  a  mixture  of  red  phosphorus,  glue,  and  powdered  glass. 


400  INORGANIC  CHEMISTRY 

The  law  imposing  the  tax  on  phosphorus  matches  (two  cents 
per  hundred  matches)  was  passed  mainly  to  protect  workmen  from 
the  disease  caused  by  breathing  fumes  of  phosphorus. 

Relation  of  Phosphorus  to  Life.  —  Phosphorus  is  essential 
to  the  growth  of  plants  and  animals.  Plants  take  phosphates 
from  the  soil  and  store  up  the  phosphorus  compounds,  espe- 
cially in  the  fruit  and  seeds.  Animals  eat  this  vegetable 
matter,  assimilate  the  phosphorus  compounds,  and  deposit 
them  in  the  bones,  brain,  and  nerve  tissue.  Bones  contain 
about  80  per  cent  of  calcium  phosphate  (Ca3(PO4)2).  Part 
of  the  combined  phosphorus  consumed  by  animals  is  re- 
jected by  them,  and  often  finds  its  way  back  into  the  soil. 

The  constant  removal  of  phosphates  by  plants  would  soon 
exhaust  the  soil.  Hence  phosphorus  is  restored  to  the  soil 
in  the  form  of  natural  or  artificial  fertilizers.  Natural  fer- 
tilizers are  (1)  stable  refuse,  which  always  contains  some 
of  the  phosphates  from  the  food  originally  fed  to  the  ani- 
mals; (2)  guano,  which  is  the  dried  phosphatic  and  nitrog- 
enous excrement  of  the  sea  birds  that  once  lived  in  vast 
numbers  in  Peru  and  Chile  ;  and  (3)  phosphate  slag,  which 
is  a  phosphorus  by-product  obtained  in  manufacturing  steel. 
Artificial  fertilizers  are  made  from  phosphate  rock.  This 
occurs  in  large  beds  in  South  Carolina,  Tennessee,  and 
Florida,  which  yield  about  a  million  tons  a  year.  It  consists 
of  the  hardened  remains  of  land  and  marine  animals,  and  is 
mainly  tricalcium  phosphate  (Ca3(PO4)2).  It  is  insoluble 
in  water,  and  must  be  changed  into  the  soluble  primary  cal- 
cium salt  (H4Ca(PO4)2)  before  it  can  be  easily  taken  up  by 
plants.  This  soluble  salt  is  called  "superphosphate  of  lime." 
When  phosphate  rock  is  treated  with  sulphuric  acid,  the 
changes  involved  may  be  written  thus  :  — 


Ca3(PO4)2  +  2  HaSO,  =  H4Ca(PO4)2  +  2CaSO4 

Tricalcium  "  Superphosphate        Calcium 

Phosphate  of  Lime"  Sulphate 


ARSENIC  401 

Ca3(P04)a  +  3H.,S04   =  2H3PO4    +  3CaS04 

Phosphoric 
Acid 

Ca3(P04)2  +  HsSO*      =  H2Ca2(P04)2  +  CaSO4 

Dicalcium 
Phosphate 

The  aim  is  to  convert  the  crude  phosphate  rock  into  "  super- 
phosphate," but  the  three  reactions  usually  occur.  The 
product  is  ground,  dried,  and  packed  in  bags  for  the  market. 
On  standing,  it  may  undergo  "  reversion,"  i.e.  the  "  super- 
phosphate" and  the  phosphoric  acid  may  form  insoluble 
phosphates,  thus  making  the  fertilizer  less  valuable.  Some- 
times "  superphosphate "  is  mixed  with  compounds  of  nitro- 
gen and  of  potassium  to  produce  a  complete  fertilizer. 

ARSENIC 

Occurrence. — Arsenic  is  found  free  in  nature,  but  it 
usually  occurs  combined  with  sulphur,  a  metal,  or  both. 
The  common  arsenic  ores  are  realgar  (As2S2),  orpiment 
(AsgSs),  arsenic  pyrites  or  mispickel  (FeSAs).  Arsenic 
trioxide  or  arsenolite  (As2O3)  is  also  found.  Small  quantities 
of  arsenic  occur  in  many  ores. 

Preparation  and  Properties.  —  Arsenic  is  prepared  in  the 
laboratory  by  reducing  arsenious  oxide  with  charcoal  in  a 
glass  tube;  the  arsenic  is  volatile  and  is  deposited  as  a  dark 
ring  on  the  colder  part  of  the  tube.  The  change  is  repre- 
sented thus : — 

2As203  +    6C    =     As4    +     6CO 

Arsenious        Carbon         Arsenic  Carbon 

Oxide  Monoxide 

On  a  large  scale  arsenic  is  extracted  from  its  ores  by  the 
method  just  indicated  or  by  roasting  arsenic  pyrites  (FeSAs) 
in  the  absence  of  oxygen. 


402  INORGANIC   CHEMISTRY 

Arsenic  is  a  brittle,  steel-gray  solid.  A  freshly  broken  piece 
has  a  metallic  luster,  which  disappears  slowly  in  a  moist 
atmosphere.  It  tends  to  crystallize.  The  specific  gravity 
varies  from  5.62  to  5.96.  Heated  in  the  air,  it  volatilizes 
without  melting,  and  the  vapor  has  an  odor  like  garlic.  At 
about  180°  C.  it  burns  in  the  air  with  a  bluish  flame,  forming 
white  arsenious  oxide  (As2O3). 

The  vapor  density  of  arsenic  at  about  650°  C.  corresponds 
to  the  molecular  weight  300.  Since  the  atomic  weight  is  75, 
the  molecular  formula  of  arsenic  vapor  is  As4  at  this  tempera- 
ture. At  about  1700°  C.  the  formula  is  As2. 

Metallic  arsenic  has  few  uses,  the  main  one  being  to  harden 
the  lead  which  is  made  into  shot. 

Arsenious  Oxide  or  Arsenic  Trioxide,  As2O3  or  As4O6,  is  the 

most  important  compound  of  arsenic,  and  is  often  called 
simply  ''arsenic"  or  "white  arsenic."  Small  quantities  are 
found  free  in  nature.  The  commercial  substance  is  obtained 
as  a  by-product  in  the  roasting  of  ores  containing  arsenic. 
There  are  two  common  varieties,  a  white,  granular  powder 
and  an  amorphous,  glasslike  solid.  It  is  an  odorless  white 
solid  with  a  faint  metallic  taste;  it  dissolves  only  slightly 
in  cold  water,  but  is  transformed  readily  by  hot  hydro- 
chloric acid  into  soluble  arsenic  trichloride  (AsCl3).  Arsenic 
trioxide  is  a  violent  poison.  The  antidote  is  fresh  ferric 
hydroxide,  which  may  be  quickly  made  by  adding  ammonium 
hydroxide  to  a  ferric  salt,  e.g.  ferric  chloride;  the  efficiency 
of  the  antidote  depends  upon  the  fact  that  the  ferric  hy- 
droxide (formed  by  the  interaction  of  ammonium  hydrox- 
ide and  ferric  chloride)  produces  an  insoluble  compound 
with  the  arsenic  compound.  Small  doses  of  arsenic  trioxide 
(2  to  3  grains)  are  usually  fatal,  but  by  habitual  use  the  sys- 
tem appropriates  larger  doses  without  ill  effects.  Workmen 
in  arsenic  factories  often  accidentally  swallow  with  impunity 


ARSENIC  403 

quantities  which  would  ordinarily  prove  fatal.  It  is  used 
to  a  limited  extent  in  making  pigments,  in  manufacturing 
glass  and  arsenic  compounds,  in  calico  printing,  in  preserv- 
ing skins,  and  in  preparing  certain  insect  and  vermin  poisons. 
At  ordinary  pressures  arsenic  trioxide  sublimes  without 
melting,  and  the  commercial  substance  is  purified  by  subliming 
the  impure  arsenic  dust  taken  from  the  flues  or  chambers 
connected  with  ore  furnaces.  The  vapor  density  at  about 
800°  C.  requires  the  formula  As4O6,  but  at  a  very  high  tem- 
perature the  formula  is  As203.  The  formula  is  also  As4O6 
according  to  the  boiling-point  method. 

Other  Arsenic  Compounds.  — The  native  mineral  orpi- 
ment  (As2S3)  is  used  in  making  a  yellow  paint,  and  realgar 
(As2S2)  a  red  paint.  Arsenic  forms  acids  analogous  to  the 
acids  of  phosphorus,  though  they  are  less  important.  Or- 
thoarsenic  acid  (H3AsO4)  is  a  white  deliquescent  solid  pre- 
pared by  the  interaction  of  concentrated  nitric  acid  and 
arsenic  or  arsenious  oxide.  Arsenious  acid  (H3As03)  is  known 
only  in  the  solution  of  its  corresponding  anhydride  (As2O8), 
resembling  carbonic  acid  (H2CO3)  in  this  respect.  Several 
salts  of  the  acids  of  arsenic  are  of  interest.  Sodium  arsenate 
(HNa2AsO4)  and  sodium  arsenite  (Na3A.sO3)  are  occasionally 
used  in  dyeing.  Scheele's  green  is  chiefly  copper  arsenite 
(HCuAsO8),  and  was  formerly  used  to  make  a  cheap  green 
paint  and  to  color  wall  paper.  Lead  arsenate  (Pb3(AsO4)2) 
and  Paris  green  (a  complex  compound  given  the  formula 
Cu3(AsO3)2 .  Cu(C2H3O2)2)  are  effective  insecticides  and  are 
used  to  exterminate  potato  bugs  and  other  insect  pests. 

Arsenic  trisulphide,  As2S3,  and  penta  sulphide,  As2S5,  are 
obtained  as  yellow  precipitates  by  passing  hydrogen  sul- 
phide gas  into  acid  solutions  of  arsenious  and  arsenic  com- 
pounds respectively.  The  formation  of  this  yellow  sulphide 


404  INORGANIC  CHEMISTRY 

is  one  test  for  arsenic.  These  sulphides  form  soluble  sulpho 
salts  with  alkaline  sulphides,  e.g. :  — 

.    AS2S3  +  3  (NHOaS  =  2  (NH4)3AsS3 

Arsenipus         Ammonium  Ammonium 

Sulphide  Sulphide  Sulpharsenite 

Sulpharsenates  are  formed  by  ammonium  polysulphide,  e.g.: 

As2S3  +  3  (NH4)2S  +  S2  =  2  (NH4)3AsS4 

Ammonium 
Sulpharsenate 

These  sulpho-salts  are  decomposed  by  hydrochloric  acid  into 
hydrogen  sulphide  gas  and  the  yellow  sulphides  of  arsenic. 
If  hydrogen  sulphide  is  added  to  a  neutral  or  basic  solution 
of  an  arsenious  compound,  the  trisulphide  is  formed  as  a 
colloid,  which  is  coagulated  by  adding  certain  acids  or  salts. 

Marsh's  Test  for  Arsenic.  —  Arsenic  can  be  easily  detected 
by  a  simple  method,  called  Marsh's  test.  An  apparatus  for 
generating  hydrogen  is  provided  with  a  hard  glass  horizontal 
delivery  tube,  narrowed  in  places  and  drawn  to  a  point. 
Pure  zinc,  pure  dilute  sulphuric  acid,  and  the  arsenic  solu- 
tion are  put  in  the  generator.  Hydrogen  and  gaseous 
hydrogen  arsenide  (or  arsine  (AsH3))  are  formed.  If  this 
mixture  is  lighted  at  the  end  of  the  delivery  tube,  it  burns 
with  a  bluish  flame,  and  metallic  arsenic  is  deposited  as  a 
black  metallic  coating  on  cold  porcelain  held  in  the  flame; 
or  if  the  tube  is  heated  in  front  of  a  narrow  place,  arsenic  is 
deposited  at  this  point.  This  deposit  dissolves  in  sodium 
hypochlorite  solution,  but  a  deposit  of  antimony,  similarly 
produced,  does  not  dissolve.  By  this  delicate  test  the  merest 
trace  of  arsenic  is  readily  and  positively  detected. 

ANTIMONY 

Occurrence.  —  Small  quantities  of  free  antimony  are 
found.  The  most  common  ore  is  stibnite  (SbjjSs),  which 
occurs  in  China,  Japan,  Austria-Hungary,  France,  Algeria, 


ANTIMONY  405 

Italy,    Mexico,    and    Turkey.     Deposits    in    California    and 
Nevada  are  also  utilized. 

Stibnite  was  found  in  the  fifteenth  century.  The  Latin  name  of 
antimony  is  stibium,  from  stibnite,  which  gives  the  symbol  of  the  ele- 
ment, Sb. 

Preparation  and  Properties.  —  Antimony  is  prepared  in- 
dustrially by  two  methods.  In  one  the  sulphide  is  roasted, 
and  the  oxide  thus  formed  is  reduced  with  charcoal.  Equa- 
tions representing  the  main  changes  are  — 

+    5O2    =   Sb204   +    3S02 


Antimony       Oxygen        Antimony        Sulphur 
Sulphide  Oxide  Dioxide 


The  other  method  consists  in  heating  the  sulphide  with  iron, 
the  equation  for  the  chemical  change  being  — 

Sb2S3    +    3Fe   =   2Sb    +    3FeS 

Antimony  Iron         Antimony  Iron 

Sulphide  Sulphide 

Antimony  is  a  silver-  white,  crystalline,  brittle  solid.  Its 
specific  gravity  is  6.7.  At  ordinary  temperatures  antimony 
does  not  tarnish  in  the  air,  but  when  heated,  it  burns  with  a 
bluish  flame,  forming  the  white,  powdery  antimony  trioxide 
(Sb2O3).  The  melting  point  is  about  630°  C.  Powdered 
antimony  burns  brilliantly  when  added  to  chlorine,  bromine, 
or  iodine,  owing  to  the  vigorous  and  rapid  combination  with 
these  elements.  Nitric  acid  oxidizes  it  to  Sb2O3  or  to  anti- 
monic  acid  (H8Sb04),  and  aqua  regia  converts  it  into  anti- 
mony trichloride  (SbCl3).  It  is  one  constituent  of  type 
metal  (see  Alloys  of  Lead)  and  other  alloys. 

Compounds  of  Antimony.  —  Antimony  forms  stibine 
(SbH3),  which  is  analogous  to  ammonia  (NHS)  and  arsine 
(AsH3)  ,  pyro-  and  meta-  acids,  the  oxides,  Sb203,  Sb2O4,  Sb2O5, 


406  INORGANIC  CHEMISTRY 

and  halogen  compounds.  It  also  forms  complex  compounds 
in  which  antimony  acts  as  a  metal.  Tartar  emetic,  or  potas- 
sium antimonyl  tartrate  (KSbO  .  C4H4O6),  is  a  white  solid, 
soluble  in  water;  it  is  used  as  a  medicine  and  as  a  mordant 
in  dyeing  cotton.  (The  group  SbO  is  called  antimonyl.) 
Antimony  trisulphide  (Sb2S3)  as  prepared  in  the  laboratory  is 
an  orange  red  solid;  it  is  formed  by  passing  hydrogen  sul- 
phide gas  into  a  solution  of  antimony  —  the  test  for  anti- 
mony. The  native  sulphide,  or  stibnite,  is  a  lustrous,  blue- 
gray  mineral,  often  beautifully  crystallized.  Antimony 
chloride  (SbCl3)  is  formed  by  the  action  of  chlorine  upon  the 
metal  or  by  boiling  the  metal  in  aqua  regia;  it  hydrolyzes 
readily,  forming  a  white  solid  —  antimony  oxychloride 
(SbOCl).  The  formation  of  antimony  oxychloride  is  some- 
times used  as  a  test  for  antimony,  but  the  more  common 
test  is  the  formation  of  the  orange  red  sulphide  (Sb2S3). 
The  sulphides  of  antimony  (Sb2S3  and  Sb2S5)  form  salts  with 
alkaline  sulphides,  which  behave  like  the  corresponding 
arsenic  compounds. 

BISMUTH 

Occurrence.  —  Bismuth  is  usually  found  in  the  native 
state,  though  it  is  not  abundant  nor  widely  distributed. 
The  oxide  (Bi2O3)  or  bismite,  and  the  sulphide  (BigSg)  or  bis- 
muthinite,  are  the  common  ores.  The  world's  supply  comes 
from  Saxony. 

Preparation  and  Properties.  —  Bismuth  is  prepared  from 
the  native  metal  by  melting  it  on  an  inclined  plate  and  allow- 
ing it  to  drain  away  from  the  solid  impurities.  Sometimes 
the  sulphide  is  roasted  and  the  resulting  oxide  is  reduced 
with  charcoal,  as  in  the  case  of  antimony. 

Bismuth  is  a  gray-white  metal  with  a  red  tinge.  Like 
antimony,  it  is  hard  and  brittle.  It  does  not  tarnish  in  dry 


BISMUTH  407 

air,  but  it  grows  dull  in  moist  air;  and  when  heated  in  air  it 
burns  with  a  bluish  flame,  forming  the  yellowish  oxide  (Bi2O8) 
Its  specific  gravity  is  about  9.9.  Hydrochloric  acid  does 
not  readily  attack  it ;  but  nitric  acid  converts  it  into  a  nitrate, 
hot  sulphuric  acid  into  a  sulphate,  and  aqua  regia  into  a 
chloride. 

Bismuth  melts  at  about  270°  C.  But  a  mixture  of  bismuth, 
lead,  and  tin  melts  at  a  low  temperature.  For  example, 
Newton's  metal  melts  at  95°  C.  and  Rose's  metal  at  100°  C.; 
while  Wood's  metal,  which  contains  cadmium,  melts  at  only 
66°-71°  C.  These  metallic  mixtures  are  called  fusible  metals 
or  alloys.  They  are  used  in  making  casts  of  wood  cuts;  but 
more  often  as  safety  plugs  in  steam  boilers  to  prevent  ex- 
plosions, as  a  fuse  in  electrical  apparatus  to  prevent  a  short 
circuit,  and  as  a  link  to  hold  in  place  fireproof  doors  and  the 
valves  in  the  automatic  sprinkling  apparatus  now  installed 
in  large  buildings. 

Compounds  of  Bismuth.  —  Bismuth  forms  no  compounds 
with  hydrogen.  There  are  four  oxides,  but  only  two  are 
well  known.  Bismuth  trioxide  (Bi2O8)  is  a  pale  yellow 
powder,  and  the  pentoxide  (Bi2O5)  is  a  brown  powder.  Bis- 
muth trioxide  is  used  to  fix  the  gilding  on  porcelain.  The 
trichloride  (BiCl3)  is  formed  by  treating  bismuth  with  aqua 
regia;  it  hydrolyzes  with  an  excess  of  water,  forming  basic 
bismuth  chloride  (Bi(OH)2Cl),  which  by  loss  of  water  be- 
comes bismuth  oxychloride  (BiOCl).  The  formation  of  the 
white  insoluble  oxychloride  is  a  test  for  bismuth.  Bismuth, 
being  a  metal,  forms  hydroxides  (Bi(OH)3  and  BiO  .  OH). 
Normal  bismuth  nitrate  (Bi(N03)3)  when  treated  with  hot 
water  forms  basic  bismuth  nitrate  (Bi(OH)2N03  or  BiON03). 
The  latter,  often  called  subnitrate  of  bismuth,  is  a  white 
powder,  and  is  used  as'  a  medicine  for  dyspepsia  and  as  a 
cosmetic. 


INORGANIC  CHEMISTRY 

The  Nitrogen  Family.  —  This  family  illustrates  typically 
the  relation  of  atomic  weight  to  properties,  for  these  elements 
display  a  gradual  change  from  non-metal  to  metal  as  we  pass 
from  nitrogen  (atomic  weight  14.01)  through  the  periodic 
members  to  bismuth  (atomic  weight  208.0).  Nitrogen  and 
phosphorus  are  distinctly  non-metallic,  arsenic  is  both  non- 
metallic  and  metallic,  antimony  is  increasingly  metallic, 
while  bismuth  is  a  typical  metal.  They  form  analogous 
compounds,  for  example  :  — 

NHS       NA         NA          — 

PH3  P208  PA  PC13 

AsH8  AsA  AsA  AsCl8 

SbH8  SbA  SbA  SbCl3 

BiCl3 


PROBLEMS  AND  EXERCISES 

1.  Calculate  the  percentage  composition  of  (a)  sodium  phosphate 

,  (6)  dihydrogen  sodium  phosphate  (H2NaPO4),  (c)  disodium 
phosphate  (HNa2PO4),  (d)  microcosmic  salt  (HNaNH4PO4). 

2.  How  much  phosphorus  is  needed  to  remove  the  oxygen  from  a 
liter  of  air?     (Assume  (1)  2  P  +  5  O  =  PaOs,  and  (2)  air  is  20  per  cent 
oxygen.) 

3.  How  much  phosphorus  is  there  in  a  ton  (2000  Ib.)  of  bone  ash 
if  70  per  cent  of  the  sample  is  calcium  phosphate  (Ca3(PO4)2)? 

4.  If  a  skeleton  weighs  25  Ib.  and  contains  60  per  cent  calcium 
phosphate,  how  much  phosphorus  does  it  contain  ? 

5.  What  is  the  weight  of  a  cylindrical  stick  of  ordinary  phosphorus 
10  cm.  long  and  15  mm.  in  diameter?     (SUGGESTION.  —  What  is  the 
specific  gravity  of  phosphorus?) 

6.  Calculate  the  percentage  composition  of  (a)  orpiment  (As2Ss), 
(6)  realgar  (As2S2),  (c)  white  arsenic  (AsaOs). 

7.  What  is  the  weight  of  a  piece  of  antimony  25  cm.  long,  15  cm. 
wide,  and  2  mm.  thick? 

8.  What  volume  of  chlorine  and  of  phosphorus  trichloride  is  formed 
by  the  complete  dissociation  of  20  1.  of  phosphorus  pentachloride  ? 

9.  If  18.5854  gm.  of  phosphorus  yield,  42,034  gm.  of  phosphorus 


BISMUTH  409 

pentoxide,  what  is  the  atomic  weight  of  phosphorus?     (At.  wt.  of 
oxygen  =  16.) 

10.  If  2.99091  gm.  of  antimony  combine  with  1.19495  gm.  of  sul- 
phur to  form  antimony  sulphide  (SbgSs),  what  is  the  atomic  weight 
of  antimony  if  the  atomic  weight  of  sulphur  is  32.07? 

11.  Suppose  29.5305  gm.  of  bismuth  trioxide  (Bi2O3)  yield  3.044  gm. 
of  oxygen.     If  16  is  the  atomic  weight  of  oxygen,  what  is  the  atomic 
weight  of  bismuth? 

12.  Write  the  formulas  of  the  following  compounds  by  applying 
the  principle  of  valence  (see  Chapter  XIV):    (a)  Arsenious  chloride, 
arsenic  chloride,   hydrogen  arsenide,   arsenic  iodide;     (6)  hydrogen 
antimonide,  antimonic   chloride,  antimonious  sulphide;    (c)  bismuth 
hydroxide,  bismuth    nitrate,  bismuth    sulphate;     (d)  calcium  phos- 
phide,  acid   calcium  phosphate,   phosphorous  bromide,   phosphoric 
iodide,  magnesium  phosphate  (ortho),  ammonium  magnesium  phos- 
phate. 

13.  Calculate  the  weight  of  phosphorus  theoretically  obtain- 
able from  (a)  2  metric  tons  of  phosphorite  (70  per  cent  Ca3(PO4)2), 
(6)  3  metric  tons  of  apatite  (90  per  cent  pure),  and  (c)  4  metric  tons 
of  bone  ash  (82  per  cent  Ca3(PO4)2). 

14.  How  much  (a)  metaphosphoric  acid  and  (6)  pyrophosphoric 
acid  can  be  made  from  200  gm.  of  orthophosphoric  acid  (90  per 
cent  pure)? 

15.  Calculate  the  per  cent  of  (a)  antimony  in  an  ore  containing 
37  per  cent  of  stibnite,  and  (6)  of  bismuth  in  an  ore  containing  45 
per  cent  of  bismite. 

16.  Calculate   the  simplest  formulas  corresponding  to    (a)   P 
=  25.833,    O  =  53.333,    H=  1.666,    Na=  19.166;     (6)    P  =  21.83, 
O  =  45.07,  H  =  .70,  Na=  32.39;   (c)  As  =  61,  S  =  39;  and  (d)  As 
=  70.1,  S  =  29.9. 

17.  Write  the  volumetric  equations  for  (a)  formation  of  phos- 
phorus pentoxide  from  phosphorus  vapor  and  oxygen,  and  (6)  prepa- 
ration of  phosphorus  trichloride,  and  (c)  preparation  of  phosphorus 
pentachloride. 

18.  Phosphine  (PH3)  burns  to  phosphorus  pentoxide  and  water. 
What  volume  of  oxygen  is  needed  for  800  cc.  of  phosphine? 


CHAPTER  XXIV 
Metals  and  Metallurgy 

Introduction.  —  The  elements  studied  thus  far  are  chiefly 
non-metals.  Metals,  however,  have  been  mentioned,  and 
many  of  their  properties  have  been  discussed.  In  the  present 
chapter  we  shall  review  these  properties  and  prepare  the  way 
for  a  fuller  treatment  of  the  metals. 

Metals  and  Non-metals.  —  Many  years  ago  the  chemical 
elements  were  divided  into  two  classes,  called  metals  and 
non-metals.  The  division  was  based  largely  on  the  conspicu- 
ous physical  properties  of  the  elements.  The  opaque,  lus- 
trous, more  or  less  heavy,  hard,  ductile,  malleable,  tenacious 
solids  were  called  metals.  All  gases  and  certain  solids,  such 
as  carbon,  sulphur,  phosphorus,  and  iodine,  were  called 
non-metals.  Their  chemical  properties  do  not  permit  such  a 
sharp  dividing  line,  however,  to  be  drawn  between  metals 
and  non-metals.  Some  elements  have  pronounced  properties, 
like  the  non-metal  sulphur  and  the  metal  iron;  these  are  typ- 
ical. A  few  elements  have  variable  properties;  sometimes 
they  act  as  metals  and  sometimes  as  non-metals.  Antimony 
and  arsenic  belong  to  this  border-line  class;  they  are  occa- 
sionally called  the  metalloids.  The  classification  into  metals 
and  non-metals  is  not  accurate  from  a  strictly  chemical  stand- 
point, but  it  is  serviceable.  The  use  in  common  life  of  the 
words  metallic  and  metal  seldom  leads  to  confusion. 

Although  lists  of  the  metals  and  non-metals  have  already 
been  given,  a  repetition  on  a  slightly  different  basis  is  not 
inappropriate  in  this  chapter.  The  following  is  a 

410 


METALS   AND    METALLURGY 


411 


g 

B 

,±3    "o             fl 

§  II   '  1 

£    0    K          dn 

c« 

§  s  § 

4) 

CO 

d 

a 

bfi 

1  '-i  1  1 

So 

9 

1-  i  S  I 

1 

3 

g 

JJJ 

d 

3    . 

d     g 

1 

2 

S  § 

0 

O     OQ 

0 

Nk 

CO 

^ 

§  6 

§     J3      ° 

§     o 

G      3 

bo     Pi    '^ 

.§  'i 

•5    3 

G      2 

P      co      ^J 

<     S 

2  S  ^ 

<  <j 

§    5 

d     1 

11 

d 

H    £    • 

0    S 

P 

.5 

.2 

'3 

d 

3 

I 

1 

1 

[7h  i  u 

.3    c   .2     |    c    0    s    o 

0    M    «            S    N    0    iS 

s 

d 

j3 

I  f  1  1  1  3 

•o 

o     o           o    ^2    'o 
cc    PH          O    OQ    O 

fi 

1 

1 

B 

i 

>S 

3 

g 

o 

(t> 

412  INORGANIC  CHEMISTRY 

Physical  Properties  of  Metals.  —  The  physical  properties 
of  most  metals  are  familiar.  The  properties  of  individual 
metals  vary  somewhat,  depending  upon  the  method  of  manu- 
facture and  the  temperature.  Until  recently  some  metals 
were  known  only  as  powders,  but  the  electrothermal  and 
aluminothermal  processes  now  enable  us  to  prepare  most 
of  them  in  coherent  masses.  All  metals  have  a  metallic 
luster,  i.e.  the  marked  property  of  reflecting  light  from 
their  polished  or  untarnished  surfaces.  All  are  opaque 
except  in  very  thin  films.  The  color  of  many  is  white, 
though  the  tint  varies.  Thus,  in  a  compact  state  silver, 
sodium,  aluminium,  mercury,  magnesium,  iron,  and  tin 
are  nearly  pure  white,  while  bismuth  is  reddish  white  and 
potassium  is  bluish.  Copper  is  the  only  red  metal  and  gold 
the  only  yellow  one  which  is  an  element.  When  powdered, 
several  metals  are  dark,  some  even  black.  Most  metals 
are  malleable  and  ductile,  i.e.  they  can  be  hammered  or  rolled 
into  sheets  and  drawn  into  wire.  Gold,  copper,  silver,  iron, 
platinum,  and  aluminium  possess  both  these  properties  to 
a  marked  degree;  while  lead,  tin,  and  zinc  are  very  malle- 
able though  not  especially  ductile.  Antimony  and  bismuth 
are  brittle.  The  hardness  of  metals  varies.  At  the  ordinary 
temperature  mercury  is  a  liquid,  sodium  and  lead  can  be 
cut  easily  with  a  knife,  and  so  on  through  the  list  up  to 
iridium,  which  is  as  hard  as  steel.  In  specific  gravity, 
which  was  once  thought  to  be  very  high,  the  metals  range 
between  lithium,  which  has  the  specific  gravity  .534,  and 
osmium,  which  has  the  specific  gravity  22.48.  Sodium 
and  potassium  are  lighter  than  water,  while  magnesium  has 
the  specific  gravity  1.75,  and  aluminium  2.58.  Metals 
conduct  heat  and  electricity.  They  also  vary  in  this  prop- 
erty. Silver,  copper,  and  aluminium  are  the  best  con- 
ductors, and  have  therefore  many  practical  applications. 
Bismuth  is  the  poorest  conductor. 


METALS   AND   METALLURGY  413 

Chemical  Properties  of  Metals.  —  Bases  are  formed  when 
oxides  of  metals  dissolve  in  water.  On  the  other  hand 
acids  result  from  dissolving  oxides  of  non-metals.  Thus  :  - 

CaO         +       H2O       =         Ca(OH)2 

Calcium  Calcium 

Oxide  Hydroxide 

S03         +       H20       = 


Sulphur  Sulphuric 

Trioxide  Acid 

Metals,  therefore,  are  base-forming  elements.  When  com- 
pounds of  metals  are  dissolved  in  water,  the  metal  becomes 
the  positive  ion  or  cation,  and  the  solutions  have  properties 
characteristic  of  the  metal  in  the  ionic  state.  Thus,  solu- 
tions of  copper  nitrate  and  .  copper  sulphate  respond  to 
the  same  test  for  copper  because  both  contain  copper  ions 
(Cu++).  Sometimes  the  solution  contains  the  metal  as 
part  of  a  compound  ion,  e.g.  in  solutions  of  potassium 
ferrocyanide  the  iron  is  present  as  the  ion  Fe(CN)6~~~  "~. 

Occurrence  of  Metals.  .  —  Only  a  few  metals  are  found 
uncombined  or  free  in  the  earth's  crust,  and  these  are  seldom 
pure.  Of  the  six  metals  known  to  the  ancients  (gold,  copper, 
silver,  tin,  iron,  and  lead)  all  except  tin  and  lead  are  found 
free.  The  metals  which  occur  free  in  the  earth's  crust  are 
called  native,  while  their  compounds  are  called  minerals; 
the  term  mineral,  however,  is  also  applied  to  certain  in- 
organic substances  found  in  the  earth's  crust  (e.g.  sulphur, 
graphite,  silica).  Those  minerals  from  which  metals  can 
be  profitably  extracted  are  called  ores;  sometimes  the 
term  ore  is  also  applied  to  a  rock  containing  native  metals, 
e.g.  gold  ore  or  copper  ore.  The  most  abundant  classes  of 
ores  are  oxides,  sulphides,  carbonates,  sulphates,  and 
hydroxides.  Many  ores  contain  arsenic.  Some  ores  are 
very  complex. 


414  INORGANIC  CHEMISTRY 

Preparation  of  Metals.  —  The  series  of  operations  by  which 
useful  metals  are  extracted  from  their  ores  is  called  met- 
allurgy. It  includes  preliminary  treatment,  smelting,  elec- 
trolysis, refining,  and  other  necessary  operations. 

The  object  of  the  preliminary  treatment  is  to  prepare 
the  ore  for  smelting  or  for  a  similar  operation.  The  ore 
as  it  comes  from  the  mine  is  usually  mixed  with  earthy 
matter  or  rock  called  gangue.  This  impurity  is  removed 
by  mechanical  or  chemical  processes,  sometimes  by  both. 
The  mechanical  process  illustrates  one  kind  of  preliminary 
treatment.  The  ore  is  first  crushed  in  a  stamp  mill.  This 
is  a  huge,  heavy  mortar  and  pestle.  The  pestle  falls  re- 
peatedly upon  the  ore,  which  is  slowly  fed  into  the  mortar 
or  die.  A  current  of  water -(or  air)  forces  the  fine  particles 
out  of  the  mortar  through  a  sieve.  The  lighter  particles 
of  the  impurities  are  washed  away,  and  the  metallic  grains 
are  extracted  by  another  mechanical  operation,  though 
chemical  processes  are  frequently  employed,  especially  with 
inferior  ores.  This  separation  of  the  valuable  part  of  the 
ore  from  the  gangue  and  reducing  it  to  a  smaller  bulk  is 
often  called  ore  dressing  or  concentration.  Copper  is  ex- 
tracted from  Lake  Superior  ores  mainly  by  this  method  of 
preliminary  treatment. 

Gold  and  silver  ores  are  often  treated  like  copper  and  then 
extracted  from  the  slime  by  mercury.  The  latter  operation 
is  called  amalgamation.  The  most  common  method  of 
extracting  metals  from  their  ores  is  by  smelting.  The 
process  varies  with  the  kind  and  composition  of  the  ore. 
Essentially,  it  consists  in  heating  a  mixture  of  the  ore  and 
coke  (or  coal)  in  a  furnace.  The  ores  used  must,  as  a  rule, 
be  oxides.  Sulphides,  hydroxides,  and  carbonates  are  first 
roasted  or  calcined  to  convert  them  into  oxides.  The 
essential  chemical  change  in  smelting  is  a  reduction  of  the 
oxide  by  carbon.  The  carbon  and  oxygen  unite  and  pass  off 


METALS  AND   METALLURGY 


415 


as  a  gas,  leaving  the  molten  metal  at  the  bottom  of  the 
furnace.  Limestone  or  a  similar  substance  is  often  added 
to  the  mixture  as  a  flux,  i.e.  to  facilitate  the  melting  and  to 
assist  in  removing  the  impurities  as  a  glassy  substance, 
called  slag.  The  operation  is  conducted  in  different  kinds 
of  furnaces.  Iron,  for  example,  is  melted  in  a  huge  upright 
furnace  called  a  blast  furnace  (Fig.  SO),  because  a  current 
of  air  is  forced  through  the  melted  mass  to  facilitate  the 
fusion  and  chemical  changes.  In  such  a  furnace  the  fuel 
and  ore  are  in  direct  con- 
tact. When  this  is  unde- 
sirable, the  reverberatory 
furnace  is  used  (Fig.  64). 
As  the  figure  shows,  in 
this  furnace  the  flame  is 
reflected  or  reverberated 
upon  the  ore  under  treat- 
ment. In  this  kind  of 
furnace  the  ore  may  be  oxi- 
dized or  reduced  without 
coming  in  contact  with  the 
fuel.  Some  ores  demand 
special  methods,  which 
will  be  described  in  con- 


FIG.  64.  —  Reverberatory  furnace.  The 
fire  burns  on  the  grate  G  and  the  long 
flame  which  passes  over  the  bridge  E  is 
reflected  down  by  the  sloping  roof  upon 
the  contents  of  the  furnace.  Gases 
escape  through  /.  The  charge,  which 
rests  upon  the  bed  B,  does  not  come 
in  contact  with  the  fuel,  but  is  oxidized 
or  reduced  by  the  flame. 


nection  with  the  metals. 

Electrolysis  is  used  to  extract  some  metals,  especially 
aluminium  and  magnesium.  Gold,  silver,  lead,  and  copper, 
are  purified  by  electrolysis.  A  few  metals  are  extracted  by 
a  wet  process.  That  is,  the  ores  are  dissolved,  and  the 
metal  is  then  precipitated  by  adding  some  substance  or  by 
electrolysis.  Thus,  inferior  gold  ores  are  dissolved  by  treat- 
ment with  potassium  (or  sodium)  cyanide,  and  the  gold  is 
then  precipitated  by  zinc. 

Some  metals,  hitherto  rare,  are  obtained  by  reducing  their 


416  INORGANIC  CHEMISTRY 

oxides  with  carbon  in  the  electric  furnace  or  by  heating  the 
powdered  oxide  with  aluminium,  e.g.  chromium. 

Alloys  are  mixtures  or  compounds  of  two  or  more  metals. 
Some  fused  metals  mix  in  all  proportions,  while  others  seem 
to  form  definite  compounds.  The  properties  of  alloys  vary 
with  the  constituents  and  their  proportions.  Some  alloys, 
especially  those  of  copper  and  of  lead,  have  many  industrial 
uses.  Alloys  in  which  mercury  is  a  constituent  are  called 
amalgams. 

PROBLEMS  AND  EXERCISES  (REVIEW) 

1.  What  is  the  specific  gravity  of  gold,  if  a  piece  weighs  4.676 
gm.  in  air,  and  loses  0.244  gm.  when  weighed  in  water?     (NOTE.  — 
Specific  gravity  equals  the  weight  in  air  divided  by  the  loss  of 
weight  in  water.) 

2.  The  weight  of  a  liter  of  ether  vapor  at  100°  C.  and  760  mm. 
is  2.44  gm.     What  is  the  molecular  weight  of  ether? 

3.  (a)  What  is  the  atomic  weight  of  phosphorus,  if  the  specific 
heat  is  .189?     (6)  Of  potassium,  if  the  specific  heat  is  .166?     (c)  Of 
manganese,  if  the  specific  heat  is  .122? 

4.  If  a  liter  of  neon  (at  0°  C.  and  760  mm.)  weighs  .902  gm., 
what  is  the  atomic  weight  of  this  monatomic  gas? 

5.  Complete  and  balance  the  following  equations  :   (a)  Cu++  + 

Or;     (6)    Ba++  +  -    -+SO4-~ 


+  2H++C1-. 


6.  Calculate  the  weight  of  zinc  dissolved  by  100  gm.  of  a  solu- 
tion of  hydrochloric  acid  containing  20  per  cent  by  weight  of  HC1. 

7.  What  volume  of  acetylene  (standard  conditions)  will  200 
pounds  of  calcium  carbide  yield? 

8.  Calculate  the  formulas  corresponding  to    (a)  N  =  26.168, 
Cl  =  66.355,    H  =  7.476;     (6)  N  =  22.222,    O  =  76.19,    H=  1.587; 
(c)  N  =  16.47,  O  =  56.47,  Na  —  27.06.     What  is  the  name  of  each 
compound  ? 

9.  What  is  the  weight  of  sulphur  in  20  1.  of  sulphur  dioxide 
measured  at  20°  C.  and  780  mm.  pressure? 

10.  Name  the  sodium  salt  of  (a)  nitrous  and  (6)  nitric  acid. 
Name  the  corresponding  salts  of  K,  barium,  Ca,  silver,  Pb,  zinc, 
NH4,  aluminium. 


CHAPTER  XXV 

Sodium,    Potassium,  Lithium,  and   Ammonium  —  Spectrum 

Analysis 

Introduction.  —  Sodium  and  potassium,  together  with  the 
less  common  element  lithium  and  the  rare  elements  rubid- 
ium and  caesium,  form  a  natural  family  in  Group  I  of  the 
periodic  classification,  known  as  alkali  metals.  These 
elements  and  their  corresponding  compounds  resemble 
each  other  closely. 

Compounds  of  the  hypothetical  metal  ammonium  are 
conveniently  treated  in  this  chapter  because  their  chemical 
relations  are  similar. 

Sodium  and  potassium  were  discovered  by  Sir  Humphry  Davy 
in  1807  by  the  electrolysis  of  their  hydroxides.  Bunsen,  by  means  of 
the  spectroscope,  discovered  lithium  in  1855,  caesium  in  1860,  and 
rubidium  in  1861. 

SODIUM 

Occurrence.  —  Sodium  is  not  found  free.  Sodium  chloride 
and  sodium  nitrate  are  the  most  abundant  compounds. 
Many  minerals,  rocks,  marine  plants,  and  mineral  waters 
contain  combined  sodium.  About  2.5  per  cent  of  the  earth's 
crust  is  sodium. 

The  symbol  of  sodium,  Na,  is  from  the  Latin  word  natrium,  which 
in  turn  comes  from  the  Greek  word  natron,  an  old  name  of  sodium 
carbonate. 

Preparation.  —  Sodium  is  manufactured  on  a  large  scale 
by  the  electrolysis  of  fused  sodium  hydroxide.  This  method, 

417 


418 


INORGANIC  CHEMISTRY 


used  on  a  small  scale  by  Davy  in  1807  to  isolate  sodium, 
became  practicable  only  recently  and  is  known  as  the  Castner 

method.  Figure  65  is  a  sketch 
of  the  apparatus.  The  body  of 
the  steel  cylinder  rests  within  a 
heated  flue.  The  iron  cathode 
(C)  passes  up  through  the  bottom 
of  the  cylinder  into  the  fused 
sodium  hydroxide.  A  cylindrical 
collecting  pot  (P)  terminating  in 
a  wire  gauze  surrounds  the  end  of 
the  cathode.  Several  carbon  bars 
(A,  A)  dip  into  the  vessel  from 
above,  and  constitute  the  anode. 
As  the  electrolysis  proceeds,  so- 
dium and  hydrogen  are  liberated 
at  the  cathode  and  oxygen  at 


FIG.  65.  -Apparatus   for   the 

manufacture  of  sodium  by    the  anode.     The  oxygen  escapes 
°f   S°diUm    through  a  pipe  (0)  without  com- 


ing in  contact  with  the  sodium, 
while  the  sodium  and  hydrogen  collect  in  P.  The  hydrogen 
escapes  through  the  top  of  P,  while  the  sodium,  which  is 
protected  from  the  oxidizing  action  of  the  air  by  the  hydro- 
gen, is  ladled  out  at  intervals. 

Properties.  —  Sodium  is  a  silver-white  metal.  It  can  be 
easily  cut  with  a  knife  and  molded  with  the  fingers.  It  is 
light  enough  to  float  upon  water,  since  its  specific  gravity 
is  0.9712  (at  20°  C.).  Heated  in  the  air,  it  melts  at  about 
96°  C.,  and  at  a  higher  temperature  it  burns  with  a  brilliant 
yellow  flame,  forming  sodium  peroxide  (Na202).  This 
intense  yellow  color  is  characteristic  of  sodium  and  is  a 
test  for  the  element  (free  or  combined).  In  moist  air  the 
bright  surface  quickly  tarnishes,  and  sodium  as  usually  seen 


SODIUM  419 

has  a  yellow  or  gray-brown  coating.  It  is,  therefore,  kept 
under  kerosene  or  a  liquid  free  from  water. 

It  decomposes  water  at  ordinary  temperatures,  liberating 
hydrogen  and  forming  sodium  hydroxide,  thus :  — 

2  Na  +  2  H2O  =  2  NaOH    +    H2 

Sodium         Water  Sodium         Hydrogen 

Hydroxide 

If  held  in  one  place  upon  water  by  filter  paper,  enough  heat 
is  generated  to  set  fire  to  the  hydrogen,  which  burns  with  a 
yellow  flame,  owing  to  the  presence  of  volatilized  sodium. 
If  sodium  is  melted  in  chlorine,  the  two  elements  combine 
with  a  brilliant  flame,  forming  sodium  chloride.  Davy, 
in  1810,  proved  in  this  way  that  common  salt  is  really  nothing 
but  sodium  chloride.  It  combines  with  hydrogen  and  forms 
a  white  solid  called  sodium  hydride  (NaH).  If  mixed  with 
mercury,  it  •  forms  sodium  amalgam,  which  is  sometimes 
used  instead  of  sodium  itself. 

A  molecule  of  sodium  has  been  found  to  be  monatomic 
by  the  vapor  density  and  the  freezing-point  methods. 

Sodium  is  used  in  the  laboratory  to  extract  water  from 
alcohol  and  ether  and  to  prepare  organic  compounds.  Large 
quantities  are  consumed  in  the  manufacture  of  sodium 
peroxide  (Na202)  and  sodium  cyanide  (NaCN).  Its  power  to 
reduce  oxides  gives  it  limited  use  in  preparing  certain  rare 
metals,  e.g.  zirconium,  tantalum,  niobium,  and  thorium, 
though  it  is  being  replaced  by  aluminium.  (See  Thermit.) 

Sodium  Chloride,  NaCl,  is  the  most  important  compound 
of  sodium.  It  is  one  of  the  most  abundant  substances, 
and  is  familiar  under  the  name  of  salt,  common  salt,  or 
table  salt.  The  presence  of  salt  in  the  ocean,  in  lakes  and 
springs,  and  in  the  soil  is  mentioned  in  the  oldest  historical 
records.  Sodium  chloride  constitutes  about  77  per  cent 


420  INORGANIC  CHEMISTRY 

of  the  salts  found  in  sea  water  and  by  far  the  largest  part 
of  the  salt  deposits  in  the  earth's  crust. 

Preparation  of  Salt.  —  The  chief  sources  of  salt  are  sea  water, 
rock  salt  deposits,  and  brines.  (1)  In  warm  countries,  as  on  the 
shores  of  the  Mediterranean  Sea,  shallow  ponds  of  sea  water  near  the 
shore  are  evaporated  by  exposure  to  the  sun  and  wind,  and  the  salt 
is  collected.  In  some  regions  sea  water  is  first  concentrated  by  allow- 
ing it  to  trickle  over  heaps  of  brush  and  then  evaporate  to  crystalliza- 
tion in  shallow  pans.  In  cold  countries,  as  on  the  shores  of  the  White 
Sea  in  Russia,  sea  water  is  allowed  to  freeze  and  the  ice  is  removed. 
The  ice  contains  no  salt,  so  the  operation  is  repeated  until  the  remain- 
ing liquid  becomes  concentrated  enough  to  evaporate  profitably 
over  a  fire.  (2)  Deposits  of  salt  are  found  in  many  parts  of  the  globe, 
the  most  important  being  in  England,  Austria-Hungary,  and  Ger- 
many. In  these  regions  and  some  parts  of  the  United  States  the  salt 
is  mined  and  purified  like  other  minerals.  This  variety  is  coarse  and 
often  impure,  and  is  largely  used  in  curing  meat  and  preserving  hides. 
(3)  Most  of  the  salt  produced  in  the  United  States  is  obtained  from 
natural  or  artificial  brines,  i.e.  from  concentrated  solutions  of  salt. 
Artificial  brines  are  made  by  forcing  water  into  salt  deposits.  Brines 
are  obtained  in  New  York,  Michigan,  Kansas,  Ohio,  West  Virginia, 
California,  Utah,  and  Louisiana.  They  are  evaporated  in  vats  by 
the  sun's  heat  or  by  heating  in  kettles  or  pans. 

All  these  methods  give  a  product  containing  as  impurities  salts  of 
calcium  and  magnesium,  which  are  largely  removed  by  further  special 
treatment.  According  to  the  standard  established  by  the  United 
States  Department  of  Agriculture,  dry  table  or  dairy  salt  must  not 
contain  over  1.4  per  cent  of  calcium  sulphate,  .5  per  cent  of  calcium 
and  magnesium  chlorides,  and  .1  per  cent  of  matter  insoluble  in 
water.  The  dampness  of  salt  is  due  to  traces  of  magnesium  and 
calcium  chlorides.  (See  Deliquescence.) 

Properties  and  Uses  of  Salt.  —  Salt  is  rather  soluble  in 
water,  100  gm.  of  water  dissolving  about  36  gm.  of  salt  at 
0°  C.  and  40  gm.  at  100°  C.  It  crystallizes  in  cubes,  which 
often  snap  open  sharply  (i.e.  decrepitate)  when  heated, 
owing  to  the  sudden  vaporization  of  the  inclosed  water. 
This  substance  is  an  essential  ingredient  of  the  food  of  man 
and  animals.  Besides  its  universal  domestic  use,  enormous 


SODIUM  421 

quantities  are  consumed  in  the  preparation  of  many  sodium 
and  chlorine  compounds,  especially  sodium  carbonate,  hydro- 
chloric acid,  and  bleaching  powder. 

Sodium  Carbonate,  Na2C03,  is  next  to  sodium  chloride  in 
importance.  Small  quantities  of  hydrated  sodium  car- 
bonates are  found  in  Egypt,  Russia,  and  in  California  and 
Nevada.  Formerly  it  was  obtained  from  the  ashes  of  marine 
plants,  but  sodium  chloride  is  now  the  source.  The  manu- 
facture of  sodium  carbonate  is  one  of  the  most  extensive 
chemical  industries.  Two  processes  are  used,  the  Leblanc 
and  the  Solvay. 

The  Leblanc  Process  has  three  stages.  (1)  Sodium  chloride  is 
changed  into  sodium  sulphate  by  sulphuric  acid,  the  equation  for  the 
change  being  — 

2NaCl  +  H2SO4  =  Na2SO4  +  2  HC1 

Sodium        Sulphuric         Sodium         Hydrochloric 
Chloride  Acid  Sulphate  Acid 

This  operation  is  called  the  salt  cake  process;  the  impure  product, 
called  "salt  cake,"  contains  about  95  per  cent  of  sodium  sulphate. 
The  hydrochloric  acid  is  a  profitable  by-product.  (See  Hydrochloric 
Acid.)  (2)  and  (3)  The  sodium  sulphate  is  reduced  to  sodium  sul- 
phide by  heating  the  "salt  cake"  with  coal;  and  the  resulting  sodium 
sulphide  is  changed  into  sodium  carbonate  by  heating  with  limestone. 
These  two  chemical  changes,  which  are  accomplished  by  one  operation, 
are  represented  by  the  following  equations :  — 

Na2SO4   +   20   =    Na2S   +   2  CO2 

Sodium         Carbon       Sodium          Carbon 
Sulphate  Sulphide        Dioxide 

Na2S     +     CaCO3     =     Na2CO3     +     CaS 

Sodium          Limestone  Sodium  Calcium 

Sulphide  Carbonate         Sulphide 

This  operation  is  called  the  black  ash  process.  The  product  is  a  dark 
brown  or  gray  porous  mass,  and  contains,  besides  37  to  45  per  cent 
of  sodium  carbonate,  considerable  calcium  sulphide  and  other  im- 
purities. The  sodium  carbonate  is  rapidly  separated  from  the  in- 
soluble portions  of  the  "black  ash"  by  extraction  with  a  regulated 


422  INORGANIC   CHEMISTRY 

stream  of  water.  The  concentrated  solution  of  sodium  carbonate  thus 
obtained  is  evaporated  to  crystallization,  and  the  crude  crystals  are 
ignited.  This  product  is  known  as  soda  ash,  and  from  its  solution  in 
water  are  obtained  soda  crystals  or  sal  soda  (NaaCOs  .  10  H^O). 

The  Solvay  Process,  often  called  the  ammonia-soda  process,  con- 
sists in  saturating  a  cold  concentrated  solution  of  sodium  chloride  first 
with  ammonia  gas  and  then  with  carbon  dioxide  gas.  The  equation 
for  the  complete  chemical  change  is  — 

H2O   +     NaCl     +      NH3      +      CO2     =     HNaCO3     +      NH4C1 

Water  Sodium  Ammonia  Carbon          Acid  Sodium  Ammonium 

Chloride  Dioxide  Carbonate  Chloride 

The  acid  sodium  carbonate  is  nearly  insoluble  in  the  cold  ammonium 
chloride  solution,  and  therefore  separates.  It  is  changed,  by  heating, 
into  sodium  carbonate,  thus :  — 

2HNaCO3     =     Na2CO3     +     CO2     +     H2O 

Acid  Sodium  Sodium  Carbon  Water 

Carbonate  Carbonate  Dioxide 

The  liberated  carbon  dioxide  is  used  again,  and  from  the  ammonium 
chloride  the  ammonia  is  also  recovered  and  used. 

Properties  and  Uses  of  Sodium  Carbonate.  — Crystallized 
sodium  carbonate  (Na2CO3 . 10  H2O)  is  often  called  alkali 
or  soda.  It  loses  water  in  the  air,  becoming  dull  at  first 
and  finally  falling  to  a  powder,  owing  to  the  fact  that  its 
vapor  pressure  is  greater  than  the  average  vapor  pressure 
of  the  air;  this  phenomenon,  as  already  stated,  is  called 
efflorescence.  When  heated,  it  melts  in  its  water  of  crys- 
tallization, and  finally  becomes  the  anhydrous  salt  (soda 
ash,  calcined  soda,  Na2CO3).  It  is  readily  soluble  in  water, 
and  the  solution,  which  is  strongly  alkaline,  is  widely  used 
as  a  cleansing  agent,  hence  the  name  washing  soda.  A 
water  solution  of  sodium  carbonate  is  alkaline  owing  to 
hydrolysis.  Sodium  carbonate  ionizes  into  2  Na+  and 
CO3~~,  but  the  CO3-ions  are  unstable  and  form  HCO3-ions 
with  the  H-ions  from  the  slightly  dissociated  water;  this 
removal  of  H-ions  finally  leaves  in  the  solution  sufficient 
OH-ions  to  produce  an  alkaline  reaction  to  litmus. 


SODIUM  423 

Besides  its  general  use  as  a  cleansing  agent,  enormous 
quantities  of  sodium  carbonate  are  consumed  in  the  glass 
and  soap  industries  and  in  preparing  sodium  compounds. 

Sodium  Bicarbonate,  HNaCO3,  is  formed  at  one  stage  of 
the  Solvay  process  (see  above).  It  may  also  be  prepared  by 
treating  crystallized  sodium  carbonate  with  carbon  dioxide 
gas.  It  is  a  white  solid,  and  is  less  soluble  in  water  than 
the  normal  carbonate.  When  heated  or  when  mixed  with 
an  acid  or  an  acid  salt,  sodium  bicarbonate  gives  up  carbon 
dioxide.  This  property  early  led  to  its  use  in  cooking, 
and  gave  the  names  cooking  soda,  baking  soda,  or  simply 
soda. 

Sodium  bicarbonate  is  so  called  because  in  it  the  ratio  of 
the  CO8  to  the  Na  is  twice  that  in  the  normal  carbonate 
(Na2CO3).  Although  called  acid  sodium  carbonate,  a  solu- 
tion of  the  pure  salt  is  practically  neutral,  the  ions  being 
Na+  and  HCO,-. 

Sodium  bicarbonate  is  one  ingredient  of  baking  powder 
and  of  various  mixtures  (except  yeast)  used  to  raise  bread, 
cake,  and  other  food.  The  other  essential  ingredient  is 
cream  of  tartar  (acid  potassium  tartrate),  which  has  a  mild 
acid  reaction  and  thus  liberates  carbon  dioxide  from  the 
bicarbonate  (see  p.  320).  Sour  milk,  which  contains  lactic 
acid,  is  sometimes  used  in  place  of  cream  of  tartar.  When 
pastry  is  raised  with  baking  soda  and  cream  of  tartar,  the 
escaping  carbon  dioxide  puffs  up  the  dough.  Hence  baking 
soda  is  often  called  saleratus  —  the  salt  which  aerates 
(from  the  Latin  words  sal,  salt,  and  aer,  air  or  gas).  Ef- 
fervescing powders,  such  as  Seidlitz  powder,  are  mixtures 
of  sodium  bicarbonate  and  tartaric  acid  or  one  of  its  acid 
salts.  When  the  ingredients  are  put  into  water,  carbon 
dioxide  is  liberated.  Sodium  bicarbonate  is  used  as  a 
medicine  to  neutralize  an  acid  stomach.  For  example,  the 


424  INORGANIC  CHEMISTRY 

"  soda  mints  "  sometimes  taken  for  this  purpose  are  mainly 
sodium  bicarbonate. 

Sodium  Hydroxide  or  Caustic  Soda,  NaOH,  is  a  white 
deliquescent,  corrosive  solid.  It  absorbs  water  and  carbon 
dioxide  rapidly  from  the  air.  When  exposed  to  the  air,  it 
becomes  moist  at  first,  then  forms  a  concentrated  solution, 
and  ultimately  solidifies,  owing  to  transformation  into  so- 
dium carbonate.  The  deliquescence  of  sodium  hydroxide,  as 
already  stated,  is  due  to  the  fact  that  the  solution  formed 
from  the  water  deposited  on  its  surface  has  a  lower  vapor 
pressure  than  the  vapor  pressure  of  the  water  vapor  in  the 
air;  hence  the  solid  continues  to  dissolve  in  the  absorbed 
water  until  a  solution  is  produced  whose  vapor  pressure 
equals  the  pressure  of  the  water  vapor  in  the  surrounding 
air.  It  dissolves  readily  in  water  with  rise  of  temperature, 
and  the  solution  is  strongly  alkaline  owing  to  the  high  degree 
of  ionization  of  the  solute.  When  heated,  it  melts  easily, 
and  is  often  cast  into  sticks  for  use  in  the  laboratory. 

Immense  quantities  are  used  in  making  hard  soap,  paper, 
and  dyest lifts ;  in  bleaching,  and  in  refining  kerosene  oil. 

Sodium  hydroxide  is  manufactured  by  two  general  methods, 
one  electrolytic  and  the  other  purely  chemical.  In  the  elec- 
trolytic method  the  sodium  liberated  from  sodium  chloride 
interacts  with  water  and  forms  sodium  hydroxide. 

The  apparatus  used  in  one  electrolytic  process  is  shown 
in  Figure  66.  It  consists  of  a  slate  box  divided  into  one 
cathode  and  two  anode  compartments  by  partitions  extend- 
ing nearly  to  the  bottom ;  the  compartments  are  separated 
by  a  layer  of  mercury  (shown  in  black).  The  T-shaped 
anodes  (A,  A)  of  graphite  and  the  cathode  (C)  of  iron  reach 
nearly  to  the  mercury.  The  anode  compartments  contain 
sodium  chloride  solution,  while  the  cathode  compartment 
contains  sodium  hydroxide  solution;  sodium  chloride  split* 


SODIUM 


425 


tion  of  the  right  concentration  flows  slowly  and  continuously 
through  the  anode  compartments  by  means  of  the  pipes  E, 
E  (and  outlets  not  shown).  When  the  current  passes, 
chlorine  is  evolved  at  the  anodes  and  escapes  through  the 


±s 

3[ 


X   F 


ii 


i 

T™ 

1 

r 

'  —  i 

i 

i  —  * 

yr* 

x   (o\ 

FIG.  66.  —  Apparatus  for  the  manufacture  of  sodium  hydroxide  by  the 
electrolysis  of  sodium  chloride. 

pipes  D,  D ;  the  sodium  is  liberated  at  the  intermediate 
cathode  of  mercury  and  forms  an  amalgam  with  it.  By 
carefully  rocking  the  cell  on  the  device  X,  X,  the  sodium 
amalgam  alone  flows  beneath  the  partitions  into  the  cathode 
compartment,  where  the  sodium  is  liberated  at  the  iron 
cathode ;  the  sodium  at  once  reacts  with  the  water,  forming 
hydrogen  and  sodium  hydroxide.  The  hydrogen  escapes 
through  the  pipe  H,  while  the  sodium  hydroxide  solution  is 
drawn  off  (through  G)  and  replaced  by  water  (through  F). 
The  sodium  hydroxide  solution  is  evaporated  to  remove  water 
and  the  molten  mass  is  poured  into  iron  barrels  or  into  molds 
about  the  diameter  of  a  lead  pencil ;  a  flake  form  is  also  made. 
The  chlorine  is  liquefied  or  used  directly  in  manufacturing 
bleaching  powder  and  other  chlorine  compounds. 

Sodium  hydroxide  is  manufactured  to  some  extent  by  adding 
lime  to  a  boiling,  dilute  solution  of  sodium  carbonate,  but  this 
process  is  being  rapidly  superseded  by  the  electrolytic  process.  The 
essential  chemical  change  is  represented  thus :  — 


Ca(OH)2 

Calcium 
Hydroxide 


Na2C03     = 

Sodium 
Carbonate 


2  NaOH 

Sodium 
Hydroxide 


CaC03 

Calcium 
Carbonate 


The  solution  of  sodium  hydroxide  is  treated  as  described  above. 


426  INORGANIC  CHEMISTRY 


Sodium  Sulphate,  NajjSO^  is  one  of  the  products  obtained 
in  the  manufacture  of  sodium  carbonate  and  of  nitric  acid 
(see  above).  At  Stassfurt  sodium  sulphate  is  prepared  by 
cooling  to  about  0°  C.  a  mixture  of  magnesium  sulphate 
and  sodium  chloride  solutions;  the  equation  is  —  • 

MgSO4     +    2  NaCl  =  Na^04    +     MgCl2 

Magnesium  Sodium  Sodium  Magnesium 

Sulphate  Chloride          Sulphate  Chloride 

Sodium  sulphate  is  a  white  solid.  It  dissolves  readily 
in  water,  and  when  a  concentrated  solution  made  at  30°  C.  is 
cooled,  large  transparent  crystals  separate.  They  have  the 
formula  Na^CVlOHjjO  and  are  called  Glauber's  salt 
from  the  discoverer.  They  lose  water  when  exposed  to  air, 
and  the  salt  continues  to  effloresce  until  it  becomes  an 
anhydrous  powder.  The  crude  salt  is  used  in  the  glass  and 
dyeing  industries,  and  the  purified  salt  as  a  medicine. 

Sodium  Sulphite,  Na-jSOs,  is  a  white  solid  prepared  by 
passing  sulphur  dioxide  into  sodium  hydroxide  solution. 
The  crystallized  salt  has  the  formula  Na^Os  .  7  H2O.  It 
interacts  with  sulphuric  acid  as  follows  :  - 

Na2SO8     -I-     H2SO4      =      Na2SO4     +     SO2    +     H2O 

Sodium  Sodium  Sulphur 

Sulphite  Sulphate  Dioxide 

It  is  used  as  a  source  of  sulphur  dioxide  and  also  as  a  pre- 
servative. 

Sodium  Nitrate,  NaN03,  is  abundant  in  Chile,  and  is  often 
called  Chile  saltpeter.  It  is  a  white  solid,  very  soluble  in 
water,  especially  hot  water,  and  is  slightly  deliquescent. 
Large  quantities  are  used  as  a  fertilizer  (either  alone  or  mixed 
with  compounds  of  potassium  and  of  phosphorus)  and  for 
making  nitric  acid  and  potassium  nitrate. 


SODIUM  427 

The  natural  deposits  of  sodium  nitrate  are  in  a  dry  region  near 
the  coast  and  cover  over  200,000  acres.  Chile  controls  the  indus- 
try, and  exports  annually  over  a  million  tons.  The  crude  nitrate, 
which  is  called  caliche,  is  treated  with  water  and  then  purified  by 
crystallization  into  a  product  containing  94-98  per  cent  of  sodium 
nitrate.  The  final  mother  liquid  is  a  source  of  iodine.  (See  Iodine. ) 

Sodium  Nitrite,  NaNO2,  is  a  white  solid.  It  is  prepared 
by  reducing  sodium  nitrate  with  lead,  thus :  — 

NaNO3       +       Pb       =       NaNO2       +       PbO 

Sodium  Lead  Sodium  Lead 

Nitrate  Nitrite  Oxide 

It  liberates  the  oxides  of  nitrogen  (NO  and  NO2)  upon  the 
addition  of  sulphuric  acid,  and  is  used  extensively  in 
manufacturing  dyes. 

Sodium  Dioxide  or  Peroxide,  Na2O3,  is  a  yellowish  solid. 
It  is  used  to  bleach  straw  and  delicate  fabrics,  and  is  an 
oxidizing  agent.  A  fused  form  is  sold  as  "  oxone."  With 
water  it  liberates  oxygen,  according  to  the  equation  — 

2  Na202  +  2  H2O   =   O2  +   4  NaOH 

Sodium  Oxygen  Sodium 

Dioxide  Hydroxide 

Miscellaneous.  —  As  stated  above,  sodium  compounds 
impart  an  intense  yellow  color  to  a  Bunsen  flame.  Most 
sodium  compounds  dissolve  in  water  and  they  yield  a 
colorless  cation  (Na+).  The  atomic  weight  of  sodium  is 
23.00  and  its  valence  is  one. 

Sodium  cyanide  (NaCN)  is  used  to  extract  gold  from  poor 
ores.  The  sodium  phosphates,  sodium  thiosulphate,  acid 
sodium  sulphite,  sodium  silicate,  and  sodium  tetraborate 
or  borax  have  already  been  described. 


428  INORGANIC  CHEMISTRY 

POTASSIUM 

Occurrence.  —  This  metal  is  not  found  free,  but  its  com- 
pounds are  abundant.  The  common  minerals  mica  and 
feldspar  are  silicates  containing  potassium.  By  the  decay 
of  these  and  other  minerals  potassium  compounds  find  their 
way  into  the  soil,  thence  into  plants  and  animals.  Potas- 
sium salts  are  found  in  wood  ashes,  in  suint  (the  oily  sub- 
stance washed  from  sheep's  wool),  in  beet-sugar  residues, 
and  in  the  deposits  in  wine  casks.  Sea  water  and  mineral 
waters  contain  potassium  salts,  particularly  potassium 
chloride  and  potassium  sulphate.  Many  potassium  salts 
are  found  at  Stassfurt,  and  they  are  the  source  of  most 
potassium  compounds.  About  2.5  per  cent  of  the  earth's 
crust  is  potassium. 

The  Stassfurt  deposits  of  the  salts  of  potassium,  magnesium,  cal- 
cium, and  sodium  are  near  Magdeburg,  Germany.  The  deposits  are 
about  five  thousand  feet  thick  and  contain  many  salts  which  were 
deposited  ages  ago  in  beds  or  layers  during  the  slow  evaporation  of  an 
inclosed  branch  of  the  sea.  The  lowest  bed  is  an  enormous  mass  of 
more  or  less  impure  sodium  chloride  (called  in  this  case  rock  salt  or 
halite),  which  was  deposited  first.  Upon  this  rest  more  or  less  regular 
layers  of  potassium  and  magnesium  salts,  higher  still  are  calcium  salts 
and  at  the  top  is  a  thick  bed  of  sandstone.  The  different  minerals 
are  mined  separately  as  far  as  possible  and  then  separated  into  com- 
mercial products  by  an  elaborate  system  of  solution,  evaporation,  and 
crystallization.  The  most  important  Stassfurt  potassium  minerals 
are  — 

Kainite      .....  KC1,  MgSO*  .  3  H2O. 

Carnallite  .....  KC1,  MgCl2  .  6  H2O. 

Polyhalite         ....  K2SO4,  MgSO*,  2  CaSO*  .  2  H2O. 

Sylvite      .....  KC1. 

Picromerite       ....  K2SO4,MgSO4.  6  H2O. 


Many  of  the  Stassfurt  salts  belong  to  a  ciass  of  compounds 
called  double  salts,  i.e.  crystalline  compounds  of  two  or  more 
normal  salts  with  one  another.  Carnallite  is  a  double 


POTASSIUM  429 

salt,  and  its  formula  is  often  written  to  emphasize  this 
fact,  thus  KC1;  MgCl2.6H2O  instead  of  KMgCl3.6H2O. 
Dilute  aqueous  solutions  of  double  salts  contain  ions  of  the 
separate  salts  and  exhibit  no  reactions  which  indicate  com- 
bination of  ions.  There  are  other  double  salts.  (See  Alums.) 

Preparation.  —  Potassium  is  prepared  by  the  electrolysis 
of  fused  potassium  chloride.  Formerly  it  was  manufactured 
by  heating  to  a  high  temperature  a  mixture  of  potassium 
carbonate  and  carbon,  or  of  potassium  hydroxide  and  iron 
carbide. 

Properties.  —  Like  sodium,  potassium  is  a  soft,  silver- 
white  metal;  it  is  light  enough  to  float  upon  water  —  the 
specific  gravity  being  .8621  (at  20°  C.).  Its  brilliant  luster 
soon  disappears  in  air,  owing  to  rapid  oxidation.  Potassium 
as  ordinarily  seen  is  covered  with  a  grayish  coating,  and  like 
sodium  must  be  kept  under  mineral  oil.  It  melts  at  62.5° 
C.,  and  at  a  higher  temperature  burns  with  a  violet-colored 
flame.  This  color  is  characteristic  of  burning  potassium, 
and  is  a  test  for  the  metal  and  its  compounds.  Like  sodium, 
it  decomposes  water  at  ordinary  temperatures,  though  more 
energetically.  The  reaction  corresponds  to  the  equation  — 

2K       +     2H2O     =     2KOH     +     H2 

Potassium  Water  Potassium         Hydrogen 

Hydroxide 

The  heat  evolved  immediately  ignites  the  liberated  hydrogen, 
and  the  melted  potassium,  surrounded  by  a  violet  flame, 
dashes  to  and  fro  upon  the  cold  water.  Potassium  combines 
with  the  halogens  and  other  elements  more  vigorously 
than  sodium,  and  forms  analogous  compounds. 

Potassium  Chloride,  KC1,  is  found  native  as  sylvite  in 
the  Stassfurt  deposits.  It  is  also  obtained  in  large  quan- 
tities by  heating  a  concentrated  solution  of  carnallite 


430  INORGANIC  CHEMISTRY 

(KC1,  MgCl2 .  6  H20)  and  allowing  potassium  chloride  to 
crystallize  out  from  the  cool  solution.  It  is  a  white  solid 
which  crystallizes  in  cubes  and  otherwise  resembles  sodium 
chloride.  It  is  used  chiefly  to  prepare  other  potassium  salts, 
especially  the  nitrate,  hydroxide,  carbonate,  and  chlorate. 

Potassium  bromide  and  iodide,  which  are  analogous  to  the  chloride, 
have  been  described  (see  these  compounds). 

Potassium  Nitrate,  KNO3,  is  also  called  niter  and  salt- 
peter. It  is  formed  in  the  soil  near  large  cities  in  India 
and  Persia  by  the  decomposition  of  nitrogenous  organic 
matter  in  the  presence  of  potassium  salts.  (See  Nitrification.) 
It  is  manufactured  by  mixing  hot,  concentrated  solutions 
of  native  sodium  nitrate  and  potassium  chloride,  which 
interact  thus :  — • 

NaN08      +      KC1      =       KN03       +       NaCl 

Sodium  Potassium  Potassium  Sodium 

Nitrate  Chloride  Nitrate  Chloride 

The  sodium  chloride,  being  the  less  soluble  of  the  two,  pre- 
cipitates as  the  solution  cools,  and  is  removed.  By  further 
evaporation  the  potassium  nitrate  (together  with  a  little 
sodium  chloride)  crystallizes  out  as  "  niter  meal." 

Potassium  nitrate  is  a  white  solid.  It  dissolves  only  to 
a  slight  extent  in  cold  water,  but  very  freely  in  hot  water. 
It  is  not  deliquescent.  If  a  solution  is  cooled  slowly,  large 
prismatic  crystals  are  formed  which  do  not  contain  water 
of  crystallization.  Such  crystals  have  cavities  in  which 
there  is  impure  water;  this  water  cannot  be  removed  by 
drying  the  crystals.  When  heated,  these  crystals  decrepi- 
tate like  those  of  sodium  chloride.  The  taste  is  salty  and 
cooling.  It  melts  at  about  339°  C.,  and  on  further  heating 
changes  into  potassium  nitrite  (KNO2)  and  oxygen.  It 
is  stable  in  the  air  at  ordinary  temperatures,  but  at  a  high 
temperature  potassium  nitrate  gives  up  oxygen  readily, 


POTASSIUM  431 

especially  to  charcoal,  sulphur,  and  organic  matter.  This 
oxidizing  power  leads  to  its  extensive  use  in  making  gun- 
powder, fireworks,  and  matches. 

Gunpowder  is  a  mixture  of  potassium  nitrate,  charcoal,  and  sulphur. 
The  ingredients  are  first  purified,  pulverized,  and  thoroughly  mixed. 
This  mixture  is  pressed,  while  damp,  into  a  thin  sheet;  and  the  press 
cake  thus  formed  is  broken  into  small  grains,  which  are  sorted  by 
sieves.  The  grains  are  then  smoothed  or  glazed  by  rolling  them  in  a 
barrel,  again  sifted,  and  finally  dried  at  a  low  temperature.  The  pro- 
portions differ  with  the  use  of  the  powder.  Modern  black  powder 
contains  about  75  per  cent  of  potassium  nitrate,  15  of  charcoal,  and 
10  of  sulphur.  When  gunpowder  burns  in  a  closed  space,  a  large 
volume  of  gas  is  suddenly  formed.  So  enormously  is  this  gas  expanded 
by  the  heat  that  in  an  open  space  it  would  fill  several  hundred  times 
the  space  taken  by  the  powder  itself.  The  pressure  exerted  by  this 
expanding  gas  is  many  tons  to  the  square  inch.  It  is  this  pressure 
which  forces  the  charge  from  a  gun  and  tears  a  rock  to  pieces.  The 
chemical  changes  attending  the  explosion  of  confined  gunpowder  are 
complex,  as  may  be  seen  by  the  following  equation:  — 

16  KN03  +  21  C  +    5S   =  5K2CO3  +  13CO2  +  K2SO4   +   SCO 

Potassium        Carbon     Sulphur      Potassium        Carbon       Potassium         Carbon 
Nitrate  Carbonate       Dioxide        Sulphate        Monoxide 

+   2K2S2   -f    8N2 

Potassium      Nitrogen 
Sulphide 

The  equation  for  the  explosion  when  unconfined  is  much  simpler, 
thus : — 

2  KNO3  +  3  C  +  S  =  3  CO2  4-  N2  +  K2S 

Gunpowder  is  being  rapidly  replaced  by  smokeless  powders.  (See 
Cellulose.) 

Potassium  Chlorate,  KC1O3,  is  a  white,  crystalline,  lustrous 
solid.  It  tastes  like  potassium  nitrate.  It  melts  at  about 
370°  C.,  and  at  a  high  temperature  decomposes  into  oxygen 
and  potassium  chloride  as  final  products,  thus:  — 

KC108    =     KC1    +    30 

Potassium        Potassium       Oxygen 
Chlorate  Chloride 


432  INORGANIC  CHEMISTRY 

It  is  used  to  prepare  oxygen,  and  in  the  manufacture  of  dyes, 
matches,  and  fireworks.  In  the  form  of  "  chlorate  of  potash 
tablets  "  it  is  used  as  a  remedy  for  sore  throat. 

Potassium  chlorate  was  formerly  manufactured  by  passing 
chlorine  into  calcium  hydroxide  (milk  of  lime)  and  adding 
potassium  chloride  to  the  mixture.  The  simplest  equations 
for  the  complex  changes  may  be  written  thus:  — 

6  Ca(OH)2  +  6  C12  =  Ca(ClO3)2  +  5  CaCl2  +  6  H2O 

Calcium  Chlorine         Calcium  Calcium 

Hydroxide  Chlorate  Chloride 

Ca(C103)2   +   2KC1    =  2KC103  +  CaCl2 

Potassium         Potassium 
Chloride  Chlorate 

This  salt  is  now  made  by  the  electrolysis  of  a  hot,  con- 
centrated solution  of  potassium  chloride.  The  two  products 
—  chlorine  and  potassium  hydroxide  —  interact  thus  :  — 

3  C12    +     6  KOH     =     KC1O3    +     5  KC1    +     3  H2O 

Chlorine  Potassium  Potassium  Potassium  Water 

Hydroxide  Chlorate  Chloride 

Potassium  Perchlorate  (KC1O4)  is  a  white  crystalline  solid  formed  by 
heating  potassium  chlorate. 

Potassium  Carbonate,  K2COs,  is  a  white  solid.  It  is 
very  soluble  in  water  and  deliquesces  when  exposed  to  the 
air,  becoming  a  thick  liquid  at  first  and  finally  a  solid.  The 
property  of  deliquescence  is  due  to  the  fact  that  the  vapor 
pressure  of  its  saturated  solution  at  ordinary  temperatures 
is  less  than  the  average  pressure  of  the  water  vapor  in  the 
air.  The  solid  residue  is  potassium  bicarbonate,  which  is 
formed  by  the  slow  absorption  of  carbon  dioxide  from  the 
air.  A  solution  of  potassium  carbonate  has  a  marked 
alkaline  reaction.  (See  Sodium  Carbonate.)  It  was  formerly 
prepared  in  large  quantities  by  extracting  wood  ashes  with 
water  and  evaporating  the  solution  to  dry  ness;  this  process 
is  still  employed  in  some  localities.  The  crude  salt  thus 
obtained  has  long  been  called  potash  and  a  purer  product 


POTASSIUM  433 

is  known  as  pearlash.  The  term  potash  is  also  sometimes 
applied  to  potassium  hydroxide  (KOH)  and  to  potassium 
oxide  (K20).  It  is  used  in  the  manufacture  of  hard  glass, 
soft  soap,  and  potassium  compounds. 

Potassium  carbonate  is  obtained  to  some  extent  from  suint  by 
igniting  the  greasy  mass  and  extracting  the  potassium  carbonate  with 
water.  Beet-sugar  residues  also  furnish  potassium  carbonate.  After 
the  sugar  has  been  obtained  from  the  beet  sirup,  the  molasses  is 
changed  by  fermentation  into  alcohol,  which  is  distilled  off;  the 
liquid  residue  is  evaporated  to  dryness  and  ignited,  and  the  potassium 
carbonate  extracted  with  water.  Potassium  carbonate  is  prepared 
by  igniting  the  crude  cream  of  tartar  collected  from  the  deposits  in 
wine  casks,  and  for  this  reason  it  is  sometimes  called  salt  of  tartar. 
All  these  sources  emphasize  the  intimate  relation  of  potassium 
compounds  to  vegetable  and  animal  life.  Much  of  the  commercial 
potassium  carbonate  is  now  made  from  potassium  chloride  by  the 
Leblanc  process.  Another  process  is  used,  principally  in  Germany, 
owing  to  the  abundance  of  crude  potassium  salts  at  Stassfurt.  It 
consists  essentially  in  forcing  carbon  dioxide  into  a  solution  of  po- 
tassium chloride  containing  suspended  magnesium  carbonate,  de- 
composing the  complex  compound  of  potassium  and  magnesium 
(MgCOs  .  HKCOs  .  4  H2O)  with  steam,  and  evaporating  the  filtered 
solution  of  potassium  carbonate. 

The  name  potassium  comes  from  the  word  potash.  The  symbol, 
K,  is  from  kalium,  the  Latin  equivalent  of  kali,  which  is  derived  from 
an  Arabic  term  for  an  alkaline  substance. 

Potassium  Hydroxide  or  Caustic  Potash,  KOH,  is  a  white 
brittle  solid,  resembling  sodium  hydroxide.  It  absorbs 
water  and  carbon  dioxide  very  readily;  and  if  exposed  to  the 
air,  it  soon  becomes  a  thick  solution  of  potassium  carbonate. 
Like  sodium  hydroxide,  it  dissolves  in  water  with  evolution 
of  heat,  forming  a  markedly  alkaline,  caustic  solution. 
Besides  its  use  in  the  laboratory,  large  quantities  are  con- 
sumed in  making  soft  soap. 

Potassium  hydroxide  was  formerly  made  in  the  same  way 
as  sodium  hydroxide,  viz.  by  adding  lime  or  milk  of  lime  to 


434  INORGANIC   CHEMISTRY 

a  boiling  dilute  solution  of  potassium  carbonate,  the  equation 
for  the  change  being  — 

Ca(OH)2    +    K2C03      =      2KOH      +      CaCO3 

Milk  Potassium  Potassium  Calcium 

of  Lime  Carbonate  Hydroxide  Carbonate 

It  is  now  manufactured  by  the  electrolysis  of  a  solution  of 
potassium  chloride,  the  process  being  like  the  one  used  for 
sodium  hydroxide. 

Other  Potassium  Compounds.  —  Potassium  Cyanide  (KCN) 
is  a  white  solid,  very  poisonous,  very  soluble  in  water,  and 
has  an  odor  like  bitter  almonds.  (See  Cyanogen,  Chapter 
XVII.)  It  is  used  in  preparing  electroplating  solutions  and 
in  extracting  gold  from  poor  ores.  Potassium  Sulphate 
(K^O*)  is  manufactured  from  kainite,  and  is  largely  used 
as  a  fertilizer  and  in  making  potassium  compounds. 

Relation  of  Potassium  to  Life.  —  Potassium,  like  nitrogen 
and  phosphorus,  is  essential  to  the  life  of  plants  and  animals. 
The  ash  of  many  common  grains,  vegetables,  and  fruits  con- 
tains potassium  as  the  carbonate.  Potassium  salts  are 
supposed  to  assist  in  the  formation  of  starch,  just  as  phos- 
phorus is  indispensable  to  the  transformation  of  nitrogen 
compounds.  Potassium  salts  taken  from  the  soil  by  plants 
must  be  returned,  if  the  soil  is  to  be  productive.  Sometimes 
crude  kainite  is  used  extensively  as  a  fertilizer;  but  wood 
ashes,  or  the  sulphate  and  chloride,  are  often  used  to  supply 
potassium  salts.  (See  Relation  of  Phosphorus  to  Life, 
Chapter  XXIII.) 

Miscellaneous.  —  As  already  stated,  potassium  compounds 
impart  a  delicate  violet  tint  to  a  Bunsen  flame.  Most  potas- 
sium compounds  are  soluble  in  water,  and  such  solutions 
contain  colorless  potassium  ions  (K+).  The  atomic  weight 
of  potassium  is  39.10  and  the  valence  is  one. 


AMMONIUM    COMPOUNDS  435 

Lithium,  Li,  is  a  silver-white  metal  and  has  the  specific  gravity 
of  only  .534  (at  20°  C.),  being  the  lightest  of  the  metallic  elements. 
Its  compounds  are  widely  distributed  in  small  quantities  in  minerals, 
mineral  waters,  and  plants.  Lithia  water  and  citrate  of  lithium  are 
often  prescribed  as  a  remedy  for  diseases  of  the  kidneys.  Lithium 
compounds  color  the  Bunsen  flame  a  bright  red  —  a  delicate  test 
for  the  element. 

Rubidium,  Rb,  and  Caesium,  Cs,  have  properties  and  form  com- 
pounds analogous  to  those  of  potassium.. 

The  Alkali  Metals,  as  already  stated,  form  a  natural  family.  The 
properties  of  the  metals  are  quite  similar,  and  the  chemical  activity 
increases  in  passing  from  lithium  (at.  wt.  6.94)  to  caesium  (at.  wt. 
132.81).  All  decompose  water,  yielding  hydrogen  and  an  hydroxide 
of  the  metal.  The  hydroxides  are  active  bases,  and  the  familiar  ones 
long  ago  gave  the  name  alkali  to  the  family.  Analogous  compounds 
are  very  much  alike ;  indeed,  in  many  operations,  it  makes  little  dif- 
ference-whether  sodium  or  potassium  compounds  are  used,  though  the 
former  are  usually  preferred  on  account  of  their  lower  price. 


AMMONIUM  COMPOUNDS 

Introduction. — We  found  in  Chapter  XIII  that  ammonium 
(NH4)  is  a  metallic  radical,  i.e.  a  group  of  elements  which 
acts  like  an  atom  of  a  metal  in  chemical  changes.  Its  most 
familiar  compound  is  ammonium  hydroxide  (NH4OH), 
which  has  the  properties  of  a  base  and  resembles  sodium 
and  potassium  hydroxides.  Other  compounds  of  ammonium, 
especially  certain  salts,  are  analogous  to  the  corresponding 
salts  of  sodium  and  potassium.  Hence,  ammonium  com- 
pounds, except  ammonium  hydroxide,  are  appropriately 
described  in  this  chapter. 

Ammonium  Chloride,  NH4C1,  is  prepared  by  passing 
ammonia  gas  into  dilute  hydrochloric  acid,  by  mixing  am- 
monium hydroxide  and  hydrochloric  acid,  or  by  letting  the 


436  INORGANIC  CHEMISTRY 

two  gases  mingle.     The  equation  for  the  essential  reaction 
NH3       +       HC1       =      NH4C1 

Ammonia          Hydrochloric  Ammonium 

Acid  Chloride 

It  is  convenient  to  regard  this  compound  as  the  ammonium 
salt  of  hydrochloric  acid,  as  if  it  were  formed  by  replacing 
the  hydrogen  of  the  acid  by  ammonium,  just  as  sodium  forms 
sodium  chloride. 

Ammonium  chloride  is  a  white,  granular,  fibrous,  or  crys- 
talline solid,  with  a  sharp,  salty  taste.  It  dissolves  easily 
in  water,  and  in  so  doing  lowers  the  temperature  markedly. 
When  ammonium  chloride  is  heated  to  a  high  temperature 
(about  350°  C.),  it  volatilizes  and  dissociates  into  ammonia 
and  hydrogen  chloride;  these  gases  reunite  to  form  ammo- 
nium chloride  as  the  temperature  falls. 

Large  quantities  of  ammonium  chloride  are  made  at  one 
stage  of  the  manufacture  of  ammonium  hydroxide  by  passing 
the  gas  into  hydrochloric  acid.  The  crude  product  is  called 
" muriate  of  ammonia"  to  indicate  its  relation  to  muriatic 
(or  hydrochloric)  acid.  It  is  largely  used  in  charging 
Leclanche  batteries,  as  an  ingredient  of  soldering  fluids,  in 
galvanizing  iron,  and  in  textile  industries.  The  crude  salt 
is  purified  by  heating  it  gently  in  a  large  iron  or  earthen- 
ware pot,  with  a  dome-shaped  cover;  the  ammonium 
chloride  volatilizes  easily  and  then  crystallizes  in  the  pure 
state  as  a  fibrous  mass  on  the  inside  of  the  cover,  but  the 
impurities  remain  behind  in  the  vessel.  The  process  of 
vaporizing  a  solid  substance  and  then  condensing  the  vapor 
directly  into  the  solid  state  is  called  sublimation.  It  differs 
from  distillation  in  that  the  substance  does  not  pass  through 
an  intermediate  liquid  state.  The  product  of  sublimation 
is  called  a  sublimate.  Sublimed  ammonium  chloride  is 
known  as  sal  ammoniac. 


AMMONIUM    COMPOUNDS  437 

Ammonium  Sulphate,  (NH4)2SO4,  is  made  by  passing 
ammonia  gas  into  sulphuric  acid,  or  by  adding  ammonium 
hydroxide  to  the  acid,  thus:  — 

2NH4OH    +    H2S04   =    (NH4)2S04   +   2H20 

Arnmonium  Ammonium 

Hydroxide  Sulphate 

The  commercial  salt  is  a  grayish  or  yellowish  solid.  It  is 
used  as  a  component  of  fertilizers,  since  it  is  rich  in  nitrogen, 
and  in  making  ammonium  alum  and  other  ammonium  com- 
pounds. 

Ammonium  Sulphide,  (NH^gS,  is  prepared  by  passing 
hydrogen  sulphide  gas  into  ammonium  hydroxide,  thus :  - 

H2S       -f       2NH4OH        =        (NHO£      +      2H2O 

Hydrogen  Ammonium  Ammonium 

Sulphide  Hydroxide  Sulphide 

The  normal  sulphide  is  unstable  and  forms  acid  ammonium 
sulphide  (also  called  ammonium  hydrosulphide,  NH4HS). 
Hence,  a  solution  of  ammonium  sulphide,  as  usually  prepared,, 
contains  both  the  normal  and  acid  sulphide;  it  smells  of 
hydrogen  sulphide  and  ammonia,  and  if  exposed  to  the  airr 
it  ultimately  changes  into  ammonia,  sulphur,  and  water. 
A  solution  of  ammonium  sulphide  is  used  in  qualitative 
analysis  to  precipitate  as  sulphides  certain  metals  of  the 
third  group,  e.g.  nickel,  cobalt,  zinc,  and  manganese.  Am- 
monium sulphide  solutions  dissolve  sulphur,  thereby  forming 
a  solution  of  complex  sulphides  called  yellow  ammonium 
sulphide  or  ammonium  polysulphide.  The  polysulphide  is 
used  in  qualitative  analysis.  (See  Test  for  Arsenic.) 

Ammonium  Nitrate,  NH4NO3,  is  made  by  passing  ammonia 
into  nitric  acid,  or  by  allowing  ammonia  gas  and  the  vapoi 
of  nitric  acid  to  mingle,  thus:  — 

NH3      +       HN03       =       NH4NO, 

Ammonia  Nitric  Ammonium 

Acid  Nitrate 


438  INORGANIC  CHEMISTRY 

It  is  a  white  salt  which  forms  beautiful  crystals.  It  dis- 
solves easily  in  water  with  a  fall  of  temperature.  When 
gently  heated  it  decomposes  into  nitrous  oxide  (N2O)  and 
water,  and  its  chief  use  is  in  the  preparation  of  nitrous 
oxide  (see  this  compound). 

Ammonium  Carbonate,  (NH4)2C03,  as  usually  found  in 
commerce  is  a  mixture  of  acid  ammonium  carbonate 
(HNH4CO3)  and  a  related  compound.  It  is  a  white  solid ; 
on  exposure  to  air  it  loses  ammonia  and  forms  the  acid  car- 
bonate. It  is  used  to  prepare  some  kinds  of  baking  pow- 
der, to  scour  wool,  as  a  medicine,  and  to  prepare  smelling 
salts,  since  it  gives  off  ammonia  readily.  The  solution  is 
used  in  qualitative  analysis  to  precipitate  barium,  stron- 
tium, and  calcium. 

Miscellaneous.  —  Most  ammonium  compounds  dissolve 
in  water,  and  such  solutions  contain  colorless  ammonium 
ions  (NH4+).  Attempts  to  isolate  the  radical  ammonium 
have  thus  far  failed.  The  valence  of  ammonium  is  one. 

Ammonium  hydroxide  and  sodium  ammonium  phosphate 
(microcosmic  salt)  have  already  been  considered. 

SPECTRUM  ANALYSIS 

When  sodium  or  one  of  its  compounds  is  introduced  into 
a  Bunsen  flame,  a  vivid  yellow  color  is  imparted  to  the  flame; 
the  color  yielded  by  potassium  or  its  compounds  is  a  delicate 
violet.  Many  elements,  especially  the  metals,  behave 
similarily.  These  colors,  as  already  stated,  often  serve  as 
a  simple  test  for  the  element  (free  or  combined).  But  if 
only  a  minute  amount  of  the  substance  is  available  or  an 
intense  color  masks  a  faint  one,  the  test  fails  or  is  unreliable; 
in  a  few  cases,  too,  the  colors  are  much  alike.  Nevertheless, 
it  is  possible  to  detect  elements  in  a  flame  even  though  the 


SPECTRUM   ANALYSIS 


439 


color  is  faint   or  obscured.     This  is  accomplished  by  the 
spectroscope  (Fig.  67).     This  instrument  consists  essentially 


£, 


FIG.  67.  —  Diagram  of  a  spectroscope   showing  collimator,    slit  (*S),   tele- 
scope, and  lenses  (LA,  L2,  L8). 

of  three  parts,  (a)  a  tube  called  the  collimator  containing 
a  narrow  slit  at  one  end  through  which  the  light  enters,  (6) 
a  triangular  glass  prism  (placed  with  its  edges  parallel  to 
the  slit)  through  which  the  light  passes  as  it  comes  from 
the  collimator,  and  (c)  a  telescope  located  at  such  an  angle 
that  the  light  can  be  viewed  as  it  emerges  from  the  prism. 
When  ordinary  white  light,  which  consists  of  rays  of  all 
colors,  enters  the  slit  and  falls  upon  the  prism,  the  rays  of 
light  in  passing  through  the 
prism  are  bent.  That  is, 
the  numberless  rays  making 
up  the  white  light  emerge 
at  an  angle  to  the  line 
along  which  they  entered,  FlG.68._DisperSionoflightbyaPrism. 
the  red  being  bent  the  least, 

the  violet  the  most  (Fig.  68).  Consequently,  the  emergent 
light,  if  caught  upon  a  piece  of  ground  glass  or  viewed 
through  the  telescope,  is  no  longer  white,  but  is  a  continuous 
band  of  colors  arranged  like  the  familiar  colors  of  the  rain- 
bow. This  band  of  colors  consists  of  a  series  of  colored 
images  of  the  slit  and  is  called  a  spectrum.  The  production 
and  examination  of  a  spectrum  is  termed  spectrum  analysis. 


440  INORGANIC   CHEMISTRY 

White  light  contains  all  the  colors  of  the  spectrum.  But 
when  colored  light,  such  as  that  from  a  sodium  or  potassium 
flame,  is  examined  by  the  spectroscope,  instead  of  the  con- 
tinuous band  of  colors,  we  see  only  those  images  of  the  slit 
which  correspond  to  the  rays  of  light  in  the  sodium  or 
potassium  flame.  Thus,  the  sodium  flame  gives  one  image 
of  the  slit  (under  usual  conditions),  which  is  seen  as  a  bril- 
liant yellow  vertical  line,  besides  other  and  minor  ones 
against  a  black*  background;  similarly,  potassium  gives 
two  conspicuous  lines,  a  red  and  a  violet.  Each  element,  if 
heated  to  a  sufficiently  high  temperature,  has  its  own  series 
of  colored  lines,  which  is  called  its  line  spectrum.  The 
spectrum  of  some  elements  is  complex,  though  many  have 
certain  lines  which  are  so  conspicuous  that  the  element  can 
be  readily  detected  in  a  mixture. 

In  the  laboratory  the  spectroscope  is  used  to  detect  the 
presence  of  certain  elements,  more  especially  the  metals  of 
the  alkali  and  alkaline  earth  families.  If  a  small  piece  of 
the  metal  or  one  of  its  compounds  (preferably  the  chloride) 
is  put  on  a  platinum  wire  and  held  in  the  Bunsen  flame 
before  the  slit,  the  characteristic  line  spectrum  of  the  ele- 
ment can  be  readily  recognized  by  looking  into  the  telescope. 
Several  elements  can  be  detected  in  a  mixture,  for  although 
certain  lines  may  coincide  or  overlap,  other  lines  are  con- 
spicuous enough  to  reveal  the  presence  of  the  components. 
Minute  quantities  are  detected  by  the  spectroscope,  since 
the  light,  though  too  faint  to  affect  the  eye,  is  concentrated 
by  the  spectroscope  into  a  bright  line  (or  lines),  which  stands 
out  against  the  black  background.  Consequently  rare 
elements,  which  can  be  obtained  only  in  small  quantities  or 
with  great  difficulty,  are  usually  studied  spectroscopically. 
Thus,  Bunsen,  who  (with  Kirchhoff)  devised  the  improved 
spectroscope,  discovered  and  studied  the  rare  metals  rubid- 
ium and  caesium.  Mention  has  already  been  made  of  the 


PROBLEMS  AND  EXERCISES  441 

fact  that  during  the  last  few  years  the  spectroscope  has 
been  especially  serviceable  in  studying  argon  and  helium 
and  the  related  gases. 

PROBLEMS  AND  EXERCISES 

1.  How  many  pounds  of  sodium  could  be  made  (theoretically) 
from  20  metric  tons  of  sodium  hydroxide? 

2.  What  is  the  weight  of  a  cubic  meter  of  potassium  ?    (Assume 
a  cubic  centimeter  of  H2O  to  weigh  1  gm.  at  20°  C.) 

3.  How  much  sulphuric  acid  is  needed  to  convert  10  gm.  of 
sodium  chloride  into  sodium  sulphate? 

4.  What  volume  of  carbon  dioxide  (at  standard  conditions)  will 
be  formed  by  heating  72  gm.  of  sodium  bicarbonate? 

5.  What  weights  of  potassium  and  water  are  needed  to  produce 
50  1.  of  hydrogen  (standard  conditions)  ? 

6.  What  weight  of  sulphur  is  needed  to  convert  80  gm.  of  sodium 
sulphite  into  sodium  thiosulphate ?      (Equation  is  Na2SO3+ 8  = 
Na2S203.) 

7.  What  weight  of  Na^CO, .  10  H2O  will  2000  Ib.  of  NaCl  pro- 
duce? 

8.  Suppose  10  gm.  of  gunpowder,  when  exploded,  yielded  3  1.  of 
gas  measured  at  0°  C.  and  760  mm.     What  would  be  the  volume  at 
1800°?     What  pressure  would  be  exerted  if  the  volume  was  kept 
unchanged  ? 

9.  Calculate  the  simplest  formulas  corresponding  to  (a)  Na  = 
32.39,  8  =  22.54,   O  =  45.07  ;     (6)Na  =  36.5,   S  =  25.4,  O  =  38.09  ; 
(c)  Na  =  29.11,  8  =  40.50,  O  =  30.38. 

10.  Starting  with  sodium,  how  would  you  prepare  successively 
the  chloride,   sulphate,   sulphide,   carbonate,   chloride,   hydroxide, 
and  metal?     Also  from  KC1  the  following  in  succession:    KNO3, 
HNO3,   NaNO3,   HNO3,    KNO3,    KNO2?     Also   from    sodium,   its 
oxide,  hydroxide,  chloride,  acid  sulphate,  normal  sulphate? 

11.  Write  the  formulas  of  the  following  compounds  by  applying 
the  principle  of  valence  (see  Chapter  XIV) :    Sodium  chlorate,  so- 
dium acetate,  sodium  fluoride,  sodium  phosphate  (ortho),  potassium 
manganate,   acid  potassium  sulphite,   lithium  carbonate,   lithium 
chloride,  lithium  phosphate  (ortho). 

12.  Indicate  the  ions  which  are  in  dilute  aqueous  solutions  of 
the  following :  Potassium  nitrate,  ammonium  chloride,  picromerite. 


CHAPTER  XXVI 
Copper  —  Silver  —  Gold 

Introduction.  —  These  metals  are  related  and  constitute  a 
family  in  Group  I  of  the  periodic  classification,  but  they  do 
not  have  such  marked  family  characteristics  as  the  alkali 
metals.  The  metals,  as  well  as  their  alloys  and  compounds, 
have  many  domestic  and  commercial  uses. 

COPPER 

Copper  has  been  known  for  ages.  Domestic  utensils  and 
weapons  of  war  containing  copper  were  used  before  similar 
objects  of  iron.  The  Romans  obtained  copper  from  the 
island  of  Cyprus.  They  called  it  cuprium  aes  (i.e.  Cyprian 
brass),  which  finally  became  simply  cuprum.  From  cuprum 
we  obtain  the  symbol  Cu  and  the  terms  cuprous  and  cupric. 

Occurrence  of  Copper.  —  Copper,  both  free  and  combined, 
is  an  abundant  element.  Single  masses  of  native  or  metallic 
copper  weighing  many  tons  are  found  in  the  Michigan 
mines  in  the  Lake  Superior  region.  Besides  the  native 
copper,  the  most  valuable  ores  are  copper  sulphide  (chal- 
cocite  or  copper  glance,  CuaS),  copper  oxide  (cuprite  or 
ruby  ore,  Cu2O),  the  copper-iron  sulphides  (copper  pyrites 
or  chalcopyrite,  CuFeS2,  and  bornite,  CusFeS3),  and  the 
basic  carbonates  (malachite,  CuCO3.  Cu(OH)2,  and  azurite, 
2CuCO3.Cu(OH)2). 

Native  copper  comes  chiefly  from  Michigan,  the  copper- 
iron  sulphide  ores  from  Montana,  and  the  oxide  and  car- 
bonates from  Arizona. 

442 


COPPER  443 

Metallurgy  of  Copper.  —  Copper  is  extracted  from  its  ores 
by  processes  which  vary  with  the  composition  of  the  ore. 
(1)  Native  copper  ore  is  first  crushed,  then  washed  to  re- 
move rocky  impurities,  and  the  concentrated  product 
finally  melted.  (2)  The  carbonates  and  oxide  are  reduced 
by  roasting  them  with  coke  in  a  blast  furnace.  The  general 
chemical  change  is  reduction  and  may  be  represented  thus:  — 

Cu2O      +      C      =      2Cu      +      CO 

Copper  Carbon  Copper  Carbon 

Oxide  Monoxide 

(3)  The  smelting  of  copper-iron  sulphides  is  a  complicated 
process.  The  ore  is  crushed  and  washed,  and  then  roasted 
in  a  furnace.  This  operation  removes  the  adhering  rock 
and  changes  much  of  the  sulphide  into  an  oxide.  The 
roasted  mass  is  next  heated  with  coal  and  sand  in  a  shaft 
or  a  reverberatory  furnace,  whereby  the  iron  is  largely 
changed  into  a  fusible  silicate  which  runs  off  as  a  slag. 
The  remaining  "  matte,"  as  it  is  called,  contains  from  35  to 
50  per  cent  of  copper  besides  iron,  sulphur,  and  arsenic 
(as  well  as  gold  and  silver),  and  is  further  treated,  (a)  It 
is  roasted  and  melted  until  all  the  iron  and  arsenic  are  re- 
moved and  mainly  copper  sulphide  remains.  This  is  finally 
roasted  to  convert  it  partly  into  an  oxide,  and  the  mixture 
of  sulphide  and  oxide  is  again  melted;  the  sulphur  passes 
off  as  sulphur  dioxide,  and  the  copper  is  left  behind.  The 
equation  for  this  final  change  is  — 

2CuO      +      Cu2S       =       4Cu      +      SO2 

Copper  Copper  Copper  Sulphur 

Oxide  Sulphide  Dioxide 

(6)  As  a  rule  the  process  is  hastened  by  pouring  the  molten 
"  matte  "  into  a  silica-lined  converter  (see  Fig.  81)  and 
blowing  air  through  the  molten  mass.  The  sulphur  passes 
off  as  sulphur  dioxide  and  the  iron  forms  a  fusible  slag.  The 


444 


INORGANIC  CHEMISTRY 


product  is  called  blister  copper  and  is  about  98  per  cent 
pure.  It  is  cast  into  thick  plates  called  anodes  and  purified 
by  electrolysis. 

Purification  of  Copper.  —  The  anodes  are  connected  with 
the  positive  electrode  of  a  powerful  battery  or  dynamo  and 
suspended  in  a  solution  of  copper  sulphate  and  sulphuric 
acid.  Sheets  of  pure  copper  are  made  the  cathodes  and 
dip  into  the  solution  as  shown  in  Figure  69.  When  the  cur- 
_  rent  passes,  pure  copper 

leaves  the  anodes  and  be- 
comes deposited  upon  the 
cathodes;  the  impurities 
either  remain  in  solution 


0 


or 


fall   to   the  bottom   of 
the  tank  as  a  slime,  from 

FIG.  69.  —  Apparatus  for  the  prepara-    which    gold    and   Silver    are 

tion  of  pure  copper  by  electrolysis,   extracted     in     appreciable 

cttlfodi^^0^^11^'^0"6   quantities.     The  operation 

can  be  interpreted  by  prin- 
ciples already  discussed.  The  copper  ions  (Cu++)  migrate  to 
the  cathode,  where  they  lose  their  electric  charges  and  are 
deposited  as  metallic  copper  (Cu).  The  SO4-ions  (SO4~~) 
migrate  to  the  anode,  where  they  likewise  lose  their  charges, 
become  ordinary  chemical  SO4-groups,  and  unite  with  copper 
from  the  anode  to  form  CuSO4,  which,  however,  immediately 
ionizes  into  Cu"1"1"  and  SO4~~.  The  gold  and  silver  do  not 
ionize  and  the  zinc  (if  present)  remains  in  the  solvent  as 
ionic  zinc  (Zn++) .  Thus  the  solution  is  constantly  supplied 
with  ions  of  copper  and  SO4,  and  pure  copper  is  removed 
from  the  anodes  and  deposited  in  equivalent  amounts 
upon  the  cathodes.  New  anodes  and  cathodes  are  supplied 
as  needed.  Electrolytic  copper,  as  it  is  often  called,  is 
exceedingly  pure. 


COPPER  445 

Properties  of  Copper.  —  Copper  is  a  bright  metal,  distin- 
guishable from  all  others  by  its  peculiar  reddish  color.  It 
is  flexible,  hard,  and  tough;  its  malleability  and  ductility 
allow  it  to  be  drawn  out  into  wire  and  rolled  into  very  thin 
sheets.  Its  specific  gravity  varies  slightly  with  the  method 
of  treatment,  but  is  about  8.9.  The  melting  point  is  about 
1083°  C.  Copper  is  an  excellent  conductor  of  both  heat 
and  electricity.  In  dry  air,  it  turns  dull ;  and  in  moist  air 
it  gradually  becomes  coated  with  a  greenish  basic  copper 
carbonate.  Heated  in  the  air,  it  is  changed  into  the  black 
copper  oxide,  and  at  a  high  temperature  it  colors  a  flame 
emerald  green.  Copper  does  not  liberate  hydrogen  readily 
from  dilute  acids.  With  nitric  acid  it  forms  copper  nitrate 
and  oxides  of  nitrogen  (see  Oxides  of  Nitrogen);  with  hot 
concentrated  sulphuric  acid  it  yields  copper  sulphate  and 
sulphur  dioxide  (see  Sulphur  Dioxide).  Hydrochloric  acid 
has  little  effect  upon  it.  Copper  replaces  certain  metals  if 
suspended  in  solutions  of  their  compounds,  e.g.  a  clean 
copper  wire  soon  becomes  coated  with  mercury,  if  placed 
in  a  solution  of  any  mercury  compound;  on  the  other  hand, 
certain  metals,  such  as  iron,  zinc,  and  magnesium,  remove 
copper  from  its  solution,  e.g.  a  nail  or  knife  blade  soon  be- 
comes coated  with  copper  if  dipped  into  a  solution  of  any 
copper  compound.  Scrap  iron  is  sometimes  used  to  pre- 
cipitate copper  on  a  large  scale.  (See  Displacement  of 
Metals  below.) 

Uses  of  Copper.  —  Next  to  iron,  copper  is  the  most  useful 
metal.  Enormous  quantities  of  wire  are  used  in  operating 
the  telegraph,  cable,  telephone,  electric  railway,  and  electric 
light.  Sheet  copper  is  made  into  household  utensils,  boilers, 
and  stills.  Copper  bolts,  nails,  rivets  and  sheathing  are 
used  on  ships,  because  copper  is  only .  slowly  corroded  by 
moist  air  and  salt  water.  All  nations  use  copper  as  the 


446  INORGANIC  CHEMISTRY 

chief  ingredient  of  small  coins.  Electrical  apparatus  utilizes 
much  copper.  Maps,  etchings,  and  some  kinds  of  engrav- 
ings are  printed  from  copper  plates;  calico  is  printed  from 
a  copper  cylinder  upon  which  the  design  is  engraved.  Books 
are  printed  and  illustrated  from  electrotypes,  made  by  de- 
positing a  film  of  copper  upon  an  impression  of  the  type 
or  design  in  wax.  In  a  similar  way  many  objects  are  cop- 
per plated.  (See  Electrotyping  and  Electroplating,  Chap- 
ter XI.) 

Alloys  of  Copper  are  important.  Brass  is  a  bright  yellow 
alloy  containing  63  to  72  per  cent  of  copper,  the  remainder 
being  zinc.  It  is  made  by  melting  these  metals  together. 
It  can  be  drawn  into  wire,  hammered  into  any  shape,  and 
turned  in  a  lathe.  It  is  harder  than  copper,  and  melts  at  a 
lower  temperature.  On  account  of  its  durability,  elasticity, 
and  other  properties,  it  has  many  uses  for  which  copper 
and  zinc  are  not  suited.  Pinchbeck,  Muntz  metal,  Bath 
metal,  Dutch  metal  (leaf  or  "gold"),  are  varieties  of  brass. 
Muntz  metal  is  often  used  in  place  of  sheet  copper  as  sheath- 
ing for  the  bottoms  of  ships,  because  it  becomes  corroded 
very  slowly.  Typical  bronze  contains  different  proportions 
of  copper,  zinc,  and  tin;  some  antique  bronzes  contain  lead 
or  iron.  The  per  cent  of  copper  is  70  to  95,  of  zinc  1  to  25, 
of  tin  1  to  18.  The  proportions  in  the  British  bronze  coinage 
are  copper  95,  zinc  1,  tin  4.  On  account  of  its  fusibility, 
beautiful  color,  and  extreme  durability,  bronze  is  used  for 
statues,  memorial  tablets,  coins,  and  medals.  The  ancients 
made  it  into  weapons  of  war  and  household  utensils.  Cannon 
were  formerly  made  of  bronze,  but  for  this  purpose  steel  is 
now  used.  Phosphor  bronze  contains  a  small  per  cent  of 
phosphorus  and  tin,  and  manganese  bronze  about  30  per 
cent  of  manganese;  they  are  tougher  than  ordinary  bronze, 
and  are  used  to  make  steamship  propellers  and  parts  of 


COPPER  447 

machines.  Silicon  bronze  is  copper  with  traces  of  iron  and 
silicon;  its  tenacity  makes  it  especially  serviceable  for  tele- 
graph, trolley,  and  telephone  wires.  Aluminium  bronze 
contains  90  to  95  per  cent  copper,  the  rest  being  aluminium. 
It  is  a  hard,  yellow-brown,  elastic  alloy,  and  is  used  in  con- 
structing hulls  of  yachts;  its  lightness,  strength,  and  re- 
sistance to  chemicals  adapt  it  to  many  other  uses. 

Gun  metal  is  about  90  per  cent  copper  and  10  per  cent  zinc;  it  was 
formerly  used  in  making  cannon,  and  is  now  used  to  some  extent  in 
making  firearms.  Bell  metal  contains  about  75  per  cent  copper  and 
25  per  cent  zinc.  Speculum  metal  contains  about  70  per  cent  copper, 
30  per  cent  tin,  and  traces  of  zinc,  nickel,  and  iron;  it  takes  a  brilliant 
polish  and  is  used  in  optical  instruments.  The  numerous  varieties 
of  German  silver  contain  different  proportions  of  copper,  nickel,  and 
zinc.  The  per  cent  of  copper  is  50  to  60,  of  nickel  20  to  25,  and  of 
zinc  about  20.  In  color  and  luster  it  resembles  silver,  for  which  it  is 
often  substituted.  Its  power  to  conduct  electricity  is  only  slightly 
affected  by  changes  of  temperature,  hence  it  is  often  used  in  resistance 
coils.  Chinese  Pakfong  (or  paktong)  is  a  variety  of  German  silver. 
The  nickel  coins  of  Germany  and  the  United  States  contain  75  per 
cent  copper  and  25  per  cent  nickel.  Copper  is  also  a  constituent  of 
many  other  coins.  Britannia  metal  and  white  metal,  in  which  copper 
is  the  minor  constituent,  are  described  under  Alloys  of  Tin. 

Compounds  of  Copper.  —  Copper  forms  two  series  of  com- 
pounds, the  cuprous  and  the  cupric.  Thus,  there  are 
cuprous  oxide  (Cu2O)  and  cupric  oxide  (CuO),  cuprous 
chloride  (CuCl)  and  cupric  chloride  (CuCl2).  The  cuprous 
compounds  contain  a  larger  proportion  of  copper  than  the 
corresponding  cupric  compounds.  The  valence  of  copper  is 
one  in  the  cuprous  series  and  two  in  the  cupric.  Not  every 
member  of  each  series,  however,  is  important  or  even  well 
known.  Solutions  of  cupric  salts  contain  blue  cupric  ions 
(Cu++)  and  those  of  cuprous  salts  colorless  cuprous  ions 
(Cu+).  Copper  compounds  are  poisonous.  Cooking  uten- 
sils made  of  copper  should  be  used  with  care ;  vege- 


448  INORGANIC   CHEMISTRY 

tables,  acid  fruits,  and  fruit  preserves,  if  boiled  in  such 
vessels,  should  be  removed  as  soon  as  cooked.  The  vessels 
themselves  should  be  kept  bright  to  prevent  the  formation 
of  soluble  copper  salts  which  might  contaminate  the  contents. 

Cuprous  Oxide,  Cu2O,  occurs  native  as  cuprite  or  ruby 
ore.  It  can  be  obtained  as  reddish  powder  by  heating  a 
mixture  of  solutions  of  copper  sulphate,  Rochelle  salt 
(potassium  sodium  tartrate),  sodium  hydroxide,  and  glucose. 
This  oxide  colors  glass  ruby  red.  It  is  a  beautiful  mineral 
and  a  valuable  ore. 

Cupric  Oxide,  CuO,  is  a  black  solid  formed  by  heating 
copper  in  a  current  of  oxygen  or  by  heating  copper  nitrate 
or  other  cupric  salts  intensely.  It  is  reduced  to  metallic 
copper  when  heated  in  a  current  of  hydrogen  or  with  sub- 
stances containing  hydrogen  or  carbon,  thus :  — 

CuO       +      H2      =       Cu    '  +      H2O 

Copper  Hydrogen  Copper  Water 

Oxide 

This  property  has  led  to  its  use  in  determining  the  amount 
of  hydrogen  or  carbon  in  compounds.  (See  also  Gravimetric 
Composition  of  Water.) 

Copper  Sulphate  or  Cupric  Sulphate,  CuSO4,  is  the  most 
useful  compound  of  copper.  It  is  a  blue,  crystalline  solid, 
and  is  often  called  blue  vitriol  or  bluestone.  The  crystal- 
lized salt  (CuSO4 .  5  H2O)  loses  water  slowly  in  the  air;  heated 
to  240°  C.,  all  the  water  escapes,  leaving  a  whitish  powder. 
This  anhydrous  copper  sulphate,  as  it  is  often  called,  absorbs 
water  from  alcohol  and  similar  liquids ;  and  when  added  to 
water,  it  again  becomes  blue.  An  aqueous  solution  of  cop- 
per sulphate  has  an  acid  reaction.  As  already  stated,  this 
is  due  to  hydrolysis.  The  cupric  ions  (Cu++)  combine  with 


COPPER  449 

hydroxyl  ions  (from  the  slightly  dissociated  water)  to  form 
the  slightly  dissociated  cupric  hydroxide  (Cu(OH)2).  This 
removal  of  OH-ions,  although  slight,  leaves  in  the  solution 
enough  hydrogen  ions  (H+)  to  turn  blue  litmus  to  red.  (See 
Hydrolysis,  Chapter  X.) 

Copper  sulphate  is  used  in  electric  batteries,  in  making 
other  copper  salts,  in  calico  printing,  dyeing,  copper  plating, 
in  preserving  timber;  recently  very  dilute  solutions  have 
been  used  to  destroy  certain  forms  of  objectionable  organic 
matter  in  drinking  water.  It  is  a  germicide  and  is  one  in- 
gredient of  certain  mixtures  (such  as  Bordeaux  mixture) 
which  are  sprayed  upon  trees  to  destroy  fungi  and  kill  in- 
sects. 

Copper  sulphate  is  prepared  on  a  large  scale  by  treating 
copper  with  dilute  sulphuric  acid,  the  equation  for  the 
chemical  change  being  — 

2Cu     +     2H2SO4     +     02     =     2CuSO4     +     2H2O 

Copper  Sulphuric  Oxygen  Copper  Water 

Acid  Sulphate 

A  large  proportion  of  the  copper  sulphate  of  commerce  is 
obtained  as  a  by-product  in  refining  gold  and  silver  with 
sulphuric  acid  (see  below). 

Copper  Nitrate,  Cu(NO3)2,  is  a  blue,  crystalline  solid, 
formed  by  the  interaction  of  dilute  nitric  acid  and  copper 
or  copper  oxide.  It  is  a  cupric  salt.  It  is  deliquescent, 
very  soluble  in  water,  and  is  readily  decomposed  by  heat 
into  cupric  oxide  and  oxides  of  nitrogen. 

Cuprous  Sulphide,  Cu2S,  is  the  bluish  black  mineral  chalcocite. 
Cupric  sulphide,  CuS,  is  the  black  precipitate  formed  by  the  inter- 
action of  hydrogen  sulphide  and  a  cupric  salt. 

Copper  Carbonates.  —  Malachite  (CuCOs .  Cu  (OH)2)  and  azurite 
(2  CuCOs .  Cu  (OH)2)  are  basic  carbonates.  Both  occur  as  minerals, 


450  INORGANIC  CHEMISTRY 

malachite  being  bright  green  and  azurite  a  magnificent  blue.  They  are 
valuable  ores  of  copper.  Malachite  is  easily  polished  and  is  used  as 
an  ornamental  stone  for  pillars,  mosaics,  and  table  tops. 

Copper  Acetates. — Verdigris  is  copper  acetate  (Cu3  (OH)2  .  (C2H3O2)4), 
which  is  used  to  some  extent  in  making  green  paint  and  Paris  green. 
The  latter  is  copper  aceto-arsenite  and  is  used  to  exterminate  potato 
bugs  and  other  insects.  Verdigris  is  a  basic  salt. 

Copper-Ammonia  Compounds.  —  When  a  little  ammonium  hydrox- 
ide is  added  to  the  solution  of  a  copper  salt,  a  whitish,  gelatinous 
precipitate  is  formed  which  upon  the  addition  of  an  excess  of  am- 
monium hydroxide  becomes  a  deep  blue  solution.  If  cupric  sulphate 
is  used,  complex  compounds  can  be  obtained  from  this  solution,  e.g. 
Cu(NH3)4SO4.  The  blue  color  of  the  solution  is  due  to  the  com- 
plex ion  Cu(NH3)4++. 

Tests  for  Copper.  —  (1)  The  reddish  color,  peculiar  "cop- 
pery" taste,  and  green  color  imparted  to  a  Bunsen  flame 
serve  to  identify  metallic  copper.  (2)  An  excess  of  ammonium 
hydroxide  added  to  the  solution  of  a  copper  compound 
produces  a  deep  blue  solution.  (3)  A  few  drops  of  acetic 
acid  and  of  potassium  ferrocyanide  solution  added  to  a 
dilute  solution  of  a  copper  compound  produce  a  brown 
gelatinous  precipitate  of  cupric  ferrocyanide  (Cu2Fe(CN)6). 

Displacement  of  Metals.  —  The  deposition  of  metallic 
copper  when  certain  metals  are  put  into  solutions  of  cop- 
per salts  and  the  displacement  of  mercury  from  solutions 
of  its  salts  by  metallic  copper  itself  are  examples  of  a  kind 
of  chemical  change  in  which  most  metals  can  participate. 
Experiment  shows  that  familiar  metals  can  be  arranged  in  a 
series  (see  page  451)  based  on  their  displacing  power.  In 
this  series  each  free  metal  displaces  the  succeeding  metals 
from  their  solutions,  and  is  in  turn  displaced  from  solution 
by  those  metals  which  precede. 


COPPER  451 

ELECTROCHEMICAL  OR  ELECTROMOTIVE  SERIES  OF  METALS 


Magnesium  +  1.31 
Aluminium  +  1.04 
Zinc             +  0.52 
Iron              +0.19 
Cadmium     +0.16 

Cobalt        +  0.05 
Nickel        -  0.02 
Lead          -  0.12 
Tin             -0.14 
Hydrogen  -  0.24 

Copper     —  0.58 
Mercury  —0.99 
Silver       -  1.04 
Platinum  —  1.10 
Gold         -  1.7 

Metals  have  what  is  called  a  solution  pressure  which  tends 
to  cause  them  to  pass  into  solution,  that  is,  to  become  ionic. 
When  a  metal  high  in  the  series,  such  as  zinc,  is  placed  in  the 
solution  of  a  metal  lower  in  the  series,  such  as  copper,  the 
zinc  tends  to  go  into  solution  as  zinc  ions,  thereby  making 
the  solution  positive.  The  metallic  zinc  becomes  corre- 
spondingly negative  and  attracts  the  positive  copper  ions 
(already  in  the  solution),  which  are  deposited  as  metallic 
copper  on  the  zinc.  The  simplified  equation  for  the  whole 
change  is 

Zn     +     Cu++     =     Zn++     +     Cu 

Similar  changes  occur  with  other  metals  and  solutions. 

When  a  metal  is  placed  in  a  solution  of  one  of  its  own 
salts,  a  difference  in  potential  is  developed  between  the 
metal  and  the  solution ;  i.e.  the  metal  and  solution  become 
oppositely  charged.  If  the  metal,  like  zinc,  is  high  in  the 
series,  the  solution  is  positive,  but  if  well  down  on  the  list, 
e.g.  copper,  the  solution  is  negative.  The  values  in  the 
above  table  are  the  differences  in  potential  (in  volts)  between 
the  various  metals  and  normal  solutions  of  their  ions.  These 
values  may  be  used  to  calculate  the  electromotive  force  of  a 
cell.  Thus,  in  a  gravity  cell,  which  consists  essentially  of 
a  copper  electrode  in  copper  sulphate  solution  and  a  zinc 
electrode  in  zinc  sulphate  solution,  the  electromotive  force 


452  INORGANIC  CHEMISTRY 

(for  normal  solutions)  is  the  difference  between  the  electrode 
potentials,  or  +  0.52  -  (-  0.58)  =  1.10.  The  solutions  in 
the  ordinary  cell  are  not  normal,  but  the  calculated  and 
actual  values  of  the  electromotive  force  are  nearly  the  same. 
Hydrogen  is  not  a  metal  in  the  common  acceptance  of 
this  term,  but  hydrogen  ions  (H+)  are  positive,  so  hydrogen 
is  usually  included  in  the  electrochemical  series  of  the  metals. 
Its  position  is  interesting.  Metals  that  precede  hydrogen 
displace  it  from  most  acids,  while  those  that  follow  do  so 
rarely,  if  ever.  That  is,  hydrogen  is  displaced  from  its  solu- 
tions only  by  metals  having  a  greater  solution  pressure. 

SILVER 

Silver  is  one  of  the  oldest  metals.  For  ages  it  has  been 
used  in  the  form  of  ornaments,  costly  vessels,  and  coins. 
It  is  a  noble  metal,  i.e.  one  which  does  not  oxidize  readily 
in  the  air.  The  Latin  name  of  silver  is  argentum,  from 
which  the  symbol  Ag  is  derived.  The  alchemists  called  it 
luna  on  account  of  its  silvery  or  "moonlike"  appearance, 
and  its  alchemistic  symbol  was  a  crescent. 

Occurrence  of  Silver.  —  Native  silver  is  found  in  Arizona, 
Mexico,  Norway;  also  in  South  America  and  Australia. 
The  chief  ores  are  the  sulphides.  The  simple  sulphide 
(silver  glance  or  argentite,  AgsS)  is  the  richest  ore,  and  is 
found  alone  in  many  localities  in  the  United  States;  it  also 
occurs  mingled  with  sulphides  of  lead,  copper,  antimony, 
or  arsenic.  These  complex  sulphides  are  found  in  Mexico, 
Peru,  Bolivia,  Chile,  and  in  Idaho.  Small  quantities  of 
native  silver  chloride  (horn  silver,  AgCl)  are  also  found; 
it  resembles  wax  or  horn  in  softness  and  color.  Alloys  oi 
silver  with  gold,  mercury,  and  copper  are  found;  average 
California  gold  contains  about  12  per  cent  silver.  Many 
ores  contain  silver,  especially  those  of  lead;  and  this  argen- 


SILVER  453 

tiferous  (or  silver-bearing)  lead  is  one  of  the  chief  sources  of 
silver. 

Metallurgy  of  Silver.  —  Silver  is  extracted  from  its  ores 
by  two  principal  processes.  (1)  In  the  amalgamation  process 
the  powdered  ore  is  first  changed  into  silver  chloride  by 
roasting  (or  simply  mixing)  it  with  sodium  chloride.  The 
mass  is  then  reduced  to  silver  by  agitation  with  water  and 
iron  (or  an  iron  compound);  the  simplest  equation  for  this 
reaction  is :  — 

2AgCl     +     Fe     =     2Ag     +     FeCl2 

Silver  Iron  Silver  Iron 

Chloride  Chloride 

The  silver  is  removed  by  adding  mercury,  which  forms  an 
amalgam  with  the  silver,  but  not  with  the  other  substances. 
When  the  amalgam  is  heated,  the  mercury  distills  off,  and 
the  silver  —  with  some  gold  —  remains  behind.  (2)  Silver 
is  extracted  from  lead  ores  by  the  Parkes  process.  After 
the  sulphur,  arsenic,  and  other  impurities  have  been  re- 
moved from  the  lead  ores,  the  final  product  is  a  mixture  of 
lead,  silver,  and  gold.  This  is  melted  and  thoroughly 
mixed  with  zinc.  As  the  mixture  cools,  an  alloy  of  silver, 
gold,  zinc,  and  a  little  lead  rises  to  the  top,  solidifies,  and  is1 
removed.  The  remaining  lead  mixture  is  treated  again 
with  zinc.  The  alloy  of  silver,  gold,  zinc,  and  lead  is  heated 
to  volatilize  the  zinc  and  to  oxidize  (or  to  melt  away)  the 
lead.  The  silver  and  gold  are  separated  by  electrolysis  or 
by  heating  the  mixture  with  sulphuric  acid;  the  gold  is  not 
acted  upon  in  the  latter  process,  but  the  silver  forms  silver 
sulphate,  which  is  reduced  by  copper  to  metallic  silver. 

Lead  ores  containing  considerable  silver  are  sometimes 
subjected  to  cupellation  to  extract  the  silver.  The  ore  or 
alloy  is  heated  in  a  furnace  having  a  shallow  hearth  made 
of  porous,  infusible  bone  ash.  The  lead  is  thereby  changed 
into  an  oxide  (PbO,  litharge),  which  melts,  and  is  partly 


454  INORGANIC  CHEMISTRY 

driven  off  by  the  air  blast  into  pots  and  partly  absorbed  by 
the  porous  cupel.  The  silver  is  prevented  from  oxidizing 
by  the  melted  litharge,  but  toward  the  end  of  the  operation 
the  thin  film  of  litharge  bursts,  and  the  metallic  silver  ap- 
pears as  a  bright  disk,  if  the  operation  is  conducted  in  a 
furnace,  and  as  a  globule  or  button,  if  the  extraction  is 
performed  in  a  small  assay  cupel.  The  process  is  then 
stopped  and  the  silver  removed. 

Properties  of  Silver.  —  Silver  is  a  lustrous,  white  metal, 
which  takes  a  brilliant  polish.     It  is  harder  than  gold,  but 
softer  than  copper.     Like  copper,  it  is  tenacious,  ductile, 
and  malleable,  and  can  be  easily  made  into  various  shapes. 
Its  specific  gravity  is  about  10.5.     It  melts  at  about  962°  C.; 
and  fuses  readily  on  charcoal  in  the  blowpipe  flame;    it 
vaporizes  in  the  oxyhydrogen  flame  and  in  the  electric  fur- 
nace.    Molten  silver  absorbs  about  twenty  times  its  volume 
of  oxygen,  which  is  expelled  violently  when  the  silver  solidifies. 
Silver  is  an  excellent  conductor  of  heat  and  electricity,  but 
it  is  too  expensive  for  such  uses.     It  does  not  tarnish  in 
air,  unless  sulphur  compounds  are  present  (especially  hydro- 
gen sulphide),  and  then  the  familiar  brown  or  black  film  of 
silver  sulphide  is  produced.     This  blackening  is  especially 
noticed  on  silver  spoons  which  have  been  put  into  eggs  or 
mustard,  and  on  silver  coins  which  have  been  carried  in 
the  pocket,  the  sulphur  in  the  latter  case  coming  from  the 
sulphur  compounds  in  the  perspiration;    the  tarnishing  of 
household  silver  is  due  to  sulphur  compounds  in  illuminat- 
ing gas  or  the  gas  from  burning  coal.     Tarnished  silver  can 
be  quickly  cleaned  by  placing  it  in  contact  with  a  piece  of 
aluminium  in  a  hot  solution  of  sodium  bicarbonate   and 
sodium  chloride.     "  Oxidized  "  silver  is  not  oxidized,  but 
coated  with  silver  sulphide.     Silver  is  only  very  slightly 
acted  upon  by  hydrochloric  acid,  and  not  at  all  by  melted 


SILVER  455 

potassium  hydroxide,  sodium  hydroxide,  or  potassium 
nitrate.  Nitric  acid  and  hot  concentrated  sulphuric  acid 
change  it  into  the  nitrate  and  sulphate  respectively. 

Alloys  of  Silver.  —  Pure  silver  is  too  soft  for  constant 
use,  so  is  usually  hardened  by  adding  a  small  amount  of 
copper.  These  alloys  are  used  in  making  coins  and  jewelry. 
The  silver  coins  of  the  United  States  and  France  contain 
900  parts  of  silver  to  100  of  copper,  and  are  called  900  fine. 
British  silver  coins  are  925  fine ;  this  quality  is  called  "  ster- 
ling silver,"  and  from  it  much  ornamental  and  useful  silver- 
ware is  made. 

Silver  Plating.  —  Metals  cheaper  than  silver  can  be  coated 
or  plated  with  pure  silver  in  largely  the  same  way  as  copper. 
Plated  silverware  has  the  ap- 
pearance of  solid  or  pure  sil- 
ver. The  object  to  be  plated 
is  carefully  cleaned,  and  made 
the  cathode  in  an  electrolytic 

cell    Containing    a   Solution    of      FIG.  70.  —  Apparatus  for  silver  plat- 

potassium    silver    cyanide 

(KAg(CN)t).     The  anode  is 

a  plate  of  pure  silver  (Fig.  70).     Silver-coated  mirrors  are 

made  by  reducing  ammonio-silver  nitrate  with  an  organic 

compound,  e.g.  formaldehyde. 

Compounds  of  Silver.  —  The  most  important  compound 
is  silver  nitrate  (AgNO3).  It  is  a  white,  crystalline  solid, 
readily  soluble  in  water.  It  is  made  by  treating  silver  with 
nitric  acid.  The  equation  for  the  chemical  change  is  :  — 

3  Ag     +     4  HNO3     =     3  AgN03     +     NO     +     2  H20 

Silver  Nitric  Silver  Nitric  Water 

Acid  Nitrate  Oxide 


456  INORGANIC  CHEMISTRY 

Exposed  to  the  light,  it  turns  dark  if  in  contact  with  organic 
matter.  It  discolors  the  skin;  if  applied  long  enough,  it 
disintegrates  the  flesh,  and  is  often  used  by  physicians  for 
this  purpose.  Its  caustic  action  and  the  silvery  color  of 
the  metal  from  which  it  is  made  long  ago  led  to  the  name, 
lunar  caustic.  Besides  its  extensive  use  in  photography  and 
silver  plating,  silver  nitrate  is  the  essential  constituent  of 
some  indelible  inks.  Silver  chloride  (AgCl)  is  made  by  add- 
ing hydrochloric  acid  or  any  soluble  chloride  to  a  solution 
of  a  silver  compound.  Thus  formed,  it  is  a  white,  curdy 
solid,  which  turns  violet  in  the  light,  and  finally  black. 
This  action  of  light  is  more  intense  if  organic  matter  is 
present.  Silver  chloride  is  only  slightly  soluble  in  water, 
but  it  dissolves  in  ammonium  hydroxide,  in  sodium  thio- 
sulphate  solution,  and  in  potassium  (or  sodium)  cyanide 
solution;  in  each  case  a  complex  compound  is  produced. 
The  formation  and  properties  of  silver  chloride  constitute 
the  test  for  silver.  Silver  bromide  (AgBr)  and  silver  iodide 
(Agl)  are  similar  to  silver  chloride  in  their  properties.  Silver 
bromide  is  slightly  yellow  and  silver  iodide  has  a  distinct 
yellow  tinge;  compared  with  the  chloride,  both  are  less 
soluble  in  ammonium  hydroxide;  the  bromide  dissolves 
readily  and  the  iodide  only  slightly  in  sodium  thiosulphate 
solution.  Silver  bromide  and  iodide  find  extensive  applica- 
tion in  photography. 

Silver  'compounds,  if  soluble  in  water,  yield  silver  ions 
(Ag+).  Complex  ions  are  formed  with  other  solvents,  e.g. 
Ag(NH3)2+  with  an  excess  of  ammonium  hydroxide  and 
Ag(CN)2~  with  an  excess  of  potassium  cyanide  solution. 

The  atomic  weight  of  silver  is  107.88  and  the  valence  is  one. 

Photography  is  based  on  the  fact  that  silver  salts,  es- 
pecially the  bromide,  change  chemically  when  mixed  with 
organic  matter  and  exposed  to  the  light.  The  photograph 


GOLD  457 

is  taken  on  a  glass  plate,  coated  on  one  side  with  a  thin 
layer  of  gelatine,  containing  the  silver  salt.  Sometimes  a 
sheet  of  sensitized  gelatine,  called  a  film,  is  used.  The  plate 
or  film  is  placed  in  the  camera  and  exposed.  The  light, 
which  is  reflected  from  the  object  being  photographed, 
changes  the  silver  salt  in  proportion  to  its  brilliancy.  The 
plate,  however,  shows  no  change  until  it  has  been  developed. 
This  process  consists  in  treating  the  plate  with  a  reducing 
agent,  e.g.  hydroquinone,  pyrogallic  acid,  or  special  mix- 
tures. As  the  developer  acts  upon  the  silver  salt,  the  image 
appears.  This  is  really  a  deposit  of  finely  divided  silver. 
Where  the  intense  light  fell  upon  the  plate,  the  deposit  is 
heavier  than  where  little  or  no  light  fell.  Hence,  dark 
parts  of  the  object  appear  light  on  the  plate,  and  light  parts 
dark ;  and  since  the  image  is  the  reverse  of  the  object,  the 
plate  is  called  a  negative.  When  the  plate  has  been  properly 
developed,  it  still  contains  some  silver  salt  not  altered  by 
the  light ;  and  if  it  were  left  on  the  plate,  the  image  would 
be  clouded  and  finally  obliterated  by  the  light.  The  image 
is,  therefore,  fixed  by  washing  off  the  unreduced  silver  salt 
with  a  solution  of  sodium  thiosulphate  (or  "hyposulphite"). 
A  print  is  made  by  laying  sensitized  paper  upon  the  nega- 
tive and  exposing  them  to  the  sunlight,  so  that  the  light 
will  pass  through  the  negative  to  the  paper.  The  negative 
obstructs  the  light  in  proportion  to  the  thickness  of  the 
silver  deposit,  so  the  photograph  has  the  same  shading  as 
the  object.  Most  prints,  like  the  plates,  must  be  fixed. 
Sometimes  the  color  is  improved  by  toning,  i.e.  by  placing 
the  print  in  a  solution  of  gold  or  of  platinum. 

GOLD 

Gold,  like  silver,  is  one  of  the  oldest  metals.     For  ages  it 
has  been  the  most  highly  prized  of  the  metals  and  exten- 


458  INORGANIC  CHEMISTRY 

sively  used  for  personal  adornment  and  for  the  fundamental 
standard  of  value.     Chemically  it  is  a  noble  metal. 

The  Latin  name  of  gold,  aurum,  gives  the  symbol  Au. 
For  several  centuries  the  mediaeval  chemists  or  alchemists 
tried  to  produce  gold  by  the  transmutation  of  base  or  cheaper 
metals.  They  were  unsuccessful  in  their  search  for  the 
philosopher's  stone,  which  they  believed  had  power  to  effect 
this  transformation. 

Occurrence  of  Gold.  —  Gold  is  widely  distributed  in  the 
native  state,  but  not  abundantly  in  many  places.  Unlike 
copper  and  silver,  its  compounds  are  few  and  rare;  the 
only  important  ones  are  the  tellurides  (compounds  of  tel- 
lurium, e.g.  AuAgTe2)  found  in  Colorado.  Gold  is  never 
found  pure,  being  alloyed  with  silver  and  occasionally  with 
copper  or  iron.  It  is  disseminated  in  fine,  almost  invisible, 
particles  among  ores  of  other  metals,  especially  the  sul- 
phides of  copper  and  iron,  though  not  so  abundantly  as 
silver.  Much  gold  is  found  in  veins  of  quartz,  and  in  the 
sand  and  gravel  formed  from  gold-bearing  rocks.  Gold 
occurs  usually  as  dust,  scales,  or  grains,  but  occasionally 
shapeless  masses  called  "nuggets"  are  found,  varying  in 
weight  from  a  few  grams  to  many  kilograms.  The  largest 
nugget  ever  known  weighed  about  84  kg.  (184  lb.). 

The  chief  gold-producing  countries  are  the  United  States,  Aus- 
tralia, South  Africa,  and  Russia.  The  annual  production  in  the 
United  States  is  about  four  million  troy  ounces,  which  comes  largely 
from  Colorado,  California,  other  Western  states,  and  Alaska. 

Gold  Mining  and  Metallurgy.  —  Gold  was  first  obtained 
by  miners  by  washing  the  gold-bearing  sand  and  gravel  of 
a  stream  in  large  pans  or  cradles.  This  primitive  method 
was  soon  replaced  by  placer  mining  and  hydraulic  mining. 
Streams  of  water,  directed  against  the  earth  containing  the 


GOLD  459 

gold,  wash  away  the  lighter  materials,  but  leave  the  heavy 
gold  behind  as  fine  particles  called  gold  dust.  From  this 
mixture  gold  and  silver  are  extracted  by  mixing  with  mer- 
cury, or  by  passing  the  moistened  mass  over  copper  plates 
coated  with  mercury.  The  amalgam  is  then  heated,  as 
in  the  metallurgy  of  silver,  to  remove  the  mercury;  the 
residue  of  gold  and  silver  is  separated  as  described  below. 
In  vein  mining  the  gold-bearing  rock  —  usually  quartz  — 
is  crushed  and  then  washed,  and  the  gold  removed  by  mer- 
cury, as  in  placer  mining.  Low  grade  ores  and  those  con- 
taining certain  metals  cannot  be  profitably  treated  with 
mercury.  In  the  chlorination  process  the  crushed  ore  is 
roasted  and  then  revolved  in  barrels  containing  bleaching 
powder  and  sulphuric  acid;  this  operation  forms  a  soluble 
gold  chloride  (AuCl3),  from  which  the  gold  is  precipitated 
as  a  fine  powder  by  ferrous  sulphate  (or  other  reducing 
agents).  Sometimes  liquid  chlorine  is  used  in  the  chlorina- 
tion process.  In  the  cyanide  process  the  crushed  ore,  or  the 
slime  from  a  previous  extraction,  is  mixed  with  a  weak  solu- 
tion of  potassium  (or  sodium)  cyanide  in  large  vats  exposed 
to  the  air;  this  operation  changes  the  gold  into  a  soluble 
cyanide,  thus:  — 

4  Au  +  8  KCN    +    O2  +  2  H2O  =  4  KAu(CN)2  +  4  KOH 

Gold        Potassium        Oxygen        Water  Potassium  Potassium 

Cyanide  Gold  Cyanide        Hydroxide 

The  gold  is  precipitated  as  a  purple  powder  from  this  solu- 
tion by  electrolysis  or  by  treatment  with  zinc. 

Purification  of  Gold.  —  Gold  obtained  by  the  foregoing 
methods  is  impure,  —  silver,  copper,  and  lead  being  the  chief 
impurities.  The  purification  of  gold  is  accomplished  by 
chemical  or  electrolytic  processes.  By  the  old  parting 
process  known  as  quartation  an  alloy  of  gold  and  silver,  in 
which  gold  is  about  one  fourth  of  the  whole,  is  treated  with 


460  INORGANIC   CHEMISTRY 

nitric  acid;  this  operation  changes  the  silver  into  the  nitrate, 
from  which  the  pure  gold  can  be  readily  removed.  The 
metals  can  also  be  parted  by  the  cheaper  method  described 
under  silver,  viz.  by  boiling  with  concentrated  sulphuric 
acid.  By  this  treatment  the  gold,  which  is  about  one  sixth 
of  the  alloy,  is  left  as  a  brownish,  porous  mass.  It  is  washed 
and  dried,  and  then  fused  into  a  coherent  mass  with  char- 
coal and  sodium  carbonate.  These  chemical  processes  have 
been  largely  displaced  by  electrolytic  methods  of  separation. 
In  one  of  the  latter  the  anode  is  an  alloy  of  gold  and  silver 
(low  in  gold),  the  cathode  is  silver,  and  the  electrolytic  solu- 
tion is  a  dilute  silver  nitrate  solution  containing  a  small  pro- 
portion of  nitric  acid.  When  the  current  passes,  part  of  the 
silver  of  the  anode  goes  into  solution  as  the  nitrate,  while  part 
is  deposited  at  the  cathode;  the  gold  remains  at  the  anode  as 
a  fine  powder  and  is  caught  in  a  cloth  bag  which  incloses  the 
whole  anode.  In  another  electrolytic  process,  which  is 
successfully  operated  in  the  United  States  mints,  the  anode 
is  an  alloy  rich  in  gold,  the  cathode  is  pure  gold,  and  the 
electrolytic  solution  is  a  solution  of  gold  chloride  (AuCl3) 
and  hydrochloric  acid;  very  pure  gold  is  deposited  on  the 
cathode,  while  the  silver  forms  silver  chloride  around  the 
anode. 

The  purity  of  gold  is  expressed  in  carats.  Pure  gold  is  24 
carats  fine;  an  alloy  containing  22  parts  of  gold  and  2  parts 
copper  is  22  carat  gold,  while  one  containing  equal  parts  gold 
and  other  metals  is  12  carat  gold. 

Properties  of  Gold.  —  Gold  is  a  lustrous,  yellow  metal. 
It  is  about  as  soft  as  lead,  and  is  the  most  ductile  and  malle- 
able of  all  metals.  The  leaf  into  which  it  is  readily  beaten 
is  very  thin.  The  melting  point  is  1063°  C.  Gold  is  one  of 
the  heaviest  metals,  its  specific  gravity  being  about  19.  Air, 
oxygen,  hydrogen  sulphide,  water  vapor,  and  the  common 


GOLD  461 

acids  do  not  attack  it;  but  it  is  changed  into  a  gold  chloride 
(AuClo)  by  chlorine  water  and  compounds  and  mixtures 
which  liberate  chlorine,  especially  a  mixture  of  concentrated 
nitric  and  hydrochloric  acids.  This  mixture,  long  known 
as  aqua  regia,  has  already  been  discussed.  (See  Aqua  Regia.) 

Uses  of  Gold.  —  Pure  gold  is  too  soft  for  most  practical 
uses,  and  is  therefore  usually  hardened  by  adding  copper  or 
silver.  The  gold-copper  alloy  has  a  reddish  color  and  is 
often  called  "  red  gold  ";  the  gold-silver  alloy  is  paler  than 
pure  gold  and  is  sometimes  called  "  white  gold."  Gold  coins 
contain  gold  and  copper.  The  United  States  standard  gold 
coins  contain  9  parts  gold  and  1  part  copper,  while  in  England 
the  legal  standard  is  11  of  gold  to  1  of  copper.  Gold  leaf  of 
various  grades  is  used  to  ornament  books,  signs,  and  other 
objects.  Jewelers  use  gold  for  many  purposes;  such  gold 
varies  from  12  to  22  carats  in  purity.  On  account  of  its 
malleability,  feeble  chemical  action,  and  beauty,  gold  is  used 
by  dentists  for  filling  teeth. 

Compounds  of  Gold  are  readily  decomposed  by  metals, 
reducing  agents  (e.g.  ferrous  sulphate  or  hydrogen  sulphide), 
fine  solids  like  charcoal,  and  by  electrolysis.  When  gold  is 
dissolved  in  aqua  regia  and  the  excess  of  acid  removed  by 
evaporation,  the  resulting  auric  chloride  (AuCl3)  gives  with 
stannous  chloride  solution  a  beautiful  purple  precipitate, 
which  has  long  been  called  "  purple  of  Cassiu*" ;  it  is  col- 
loidal gold.  Its  formation  is  a  test  for  gold.  The  process 
of  gold  plating  is  the  same  as  silver  plating,  only  the  solution 
is  one  of  potassium  auri cyanide  (KAu(CN)4)  and  the  anode 
is  gold.  Much  cheap  jewelry  is  gold  plated. 

The  atomic  weight  of  gold  is  197.2. 

The  Copper  Family  belongs  to  Group  I  in  the  periodic 
classification,  but  it  is  not  a  typical  family.  Not  only  are 


462  INORGANIC  CHEMISTRY 

these  elements  quite  different  from  the  other  family  in  this 
group  (i.e.  the  alkali  metals),  but  the  different  members  are 
not  closely  related  to  each  other.  Indeed  copper  more  closely 
resembles  mercury  and  iron  than  it  does  silver,  while  gold 
has  properties  which  strongly  suggest  aluminium  and  iron. 
The  valence  of  copper,  as  already  stated,  is  one  in  cuprous 
compounds  and  two  in  cupric.  The  valence  of  silver  is 
usually  one.  Gold  has  the  valence  of  one  in  aurous  com- 
pounds and  three  in  auric  compounds. 

PROBLEMS  AND  EXERCISES 

1.  How  much  silver  chloride  is  formed  by  adding  10  gm.  of 
crystallized  barium  chloride  (BaCl2 .  2  H2O)  to  silver  nitrate? 

2.  How  many  pounds  of  copper  can  be  obtained  from  a  ton 
(2000  Ib.)  of  pure  chalcopyrite  (CuFeS2)? 

3.  A  flask  filled  with  water  weighed  153  gm. ;   25  gm.  of  cop- 
per were  dropped  in.     The  flask  and  contents  then  weighed  175.19 
gm.     What  is  the  sp.  gr.  of  copper? 

4.  What  weight  of  gold  will  25  gm.  of  ferrous  sulphate  precipi- 
tate from  a  solution  of   auric   chloride?      (Equation   is  AuCls  + 
3  FeSO4  =  Au  +  Fe2(SO4)3  +  FeCl3.) 

5.  What  weight  of  (a)  silver  and  (6)  gold  will  be  precipitated 
from  the  respective  solutions  by  25  gm.  of  copper? 

6.  Starting  with  silver,  how  would  you  prepare  in  succession 
silver  nitrate,  silver  chloride,  silver? 

7.  As  in  Exercise  6,  the  following  from  copper:   the  oxide, 
nitrate,  oxide,  metal,  sulphate,  metal? 

8.  Calculate  the  atomic  weight  of  copper,  silver,  or  gold  from 
the  following  :*  (a)  20.6885  gm.  of  copper  oxide  give  16.5167  gm. 
of  copper;    (6)  4.39313  gm.  of  copper  precipitated  14.9104  gm.  of 
silver ;    (c)  the  specific  heat  of  gold  is  .032. 

9.  Calculate  the  solubility  products  of  silver  chloride,  bromide, 
and  iodide  if  the  ionization  is  100  per  cent  and  the  molar  solubilities 
are  .0000094,  .00000058,  and  .000000016  respectively. 

10.  Write  the  formulas  of  the  following  compounds  by  apply- 
ing the  principle  of  valence  (see  Chapter  XIV) :  Cuprous  iodide, 
cupric  acetate,  silver  cyanide,  silver  sulphate,  silver  oxide,  aurous 
bromide,  auric  chloride,  aurous  hydroxide,  cupric  hydroxide. 


CHAPTER   XXVII 
Calcium,   Strontium,  and  Barium  —  Radium 

THESE  elements  form  a  natural  family  in  Group  II  of  the 
periodic  classification  known  as  the  alkaline  earth  metals. 

CALCIUM 

Occurrence.  —  Calcium  is  never  found  free.  Combined 
calcium  makes  up  about  3.5  per  cent  of  the  earth's  crust. 
The  most  abundant  compound  is  cal- 
cium carbonate  (CaCO3) .  Many  rocks 
are  silicates  of  calcium  and  other 
metals.  Calcium  sulphate  (CaSO4) 
is  abundant.  Calcium  compounds 
are  essential  to  the  life  of  plants 
and  animals,  being  found  in  the 
leaves  of  plants,  and  in  the  bones, 
teeth,  and  shells  of  animals.  The 
ocean,  many  rivers,  lakes,  and  springs 
contain  calcium  salts,  especially  the 
acid  carbonate  (H2Ca(CO3)2)  and  the 
sulphate. 

•E 


FIG.  70  a.  —  Apparatus 
for  preparing  calcium. 


Preparation.  —  Calcium  is  prepared 
by  the  electrolysis  of  melted  calcium 
chloride  (Fig.  70 a).  The  anode  is  a 

graphite  crucible    (A)    and  the    cathode  is  a  rod  of  iron 
(B)  which  can  be  adjusted  by  a  screw  (C)  so  that  it  will 

463 


464  INORGANIC  CHEMISTRY 

just  dip  into  the  melted  chloride.  The  lower  part  of  the 
crucible  is  cooled  by  a  current  of  water  in  EE.  When  the 
current  passes  calcium  collects  on  the  end  of  the  cathode 
and  is  gradually  removed  as  an  irregular  rod  (D)  by  slowly 
raising  the  cathode ;  the  end  of  the  calcium  rod  dips  into 
the  electrolyte  and  thus  serves  as  the  lower  end  of  the 
original  cathode. 

Properties.  —  Calcium  is  a  silvery  metal.  Its  specific 
gravity  is  about  1.55  and  its  melting  point  is  about  810°  C. 
It  tarnishes  in  the  air  and  should  be  kept  under  some  water- 
free  liquid.  When  heated  in  air  it  forms  an  oxide  (CaO) 
and  a  nitride  (Ca3N2).  It  interacts  with  water,  slowly  at 
ordinary  temperatures  and  rapidly  at  high  temperatures, 
forming  calcium  hydroxide  (Ca(OH)2)  and  hydrogen.  With 
acids  it  interacts  readily,  yielding  hydrogen  and  a  calcium 
salt. 

Calcium  Carbonate,  CaCO3.  —  The  most  abundant  form 
of  this  compound  is  limestone.  Vast  deposits  are  found  in 
many  places.  Limestone  is  a  white  or  gray  compact  solid, 
but  impurities,  especially  organic  matter  and  iron  compounds, 
produce  blue,  yellow,  reddish,  and  black  varieties.  Hard, 
crystalline  limestone  which  takes  a  good  polish  is  called 
marble.  This  form,  which  has  a  wide  range  of  color,  is  used 
as  a  building  and  an  ornamental  stone.  Calcite  is  crystal- 
lized calcium  carbonate.  It  is  almost  as  abundant  as  quartz, 
though  softer;  its  varied  color  and  crystal  form  combine  to 
make  it  an  attractive  mineral  (Fig.  71).  A  very  transparent 
variety  of  calcite  called  Iceland  spar  has  the  remarkable 
property  of  double  refraction,  i.e.  making  small  objects 
appear  double. 

Calcium  carbonate  is  not  soluble  in  water,  unless  carbon 
dioxide  is  present.  (See  also  Carbon  Dioxide  and  Acid  Cal- 


CALCIUM 


465 


cium  Carbonate.)  As  water  containing  carbon  dioxide  works 
its  way  underground  in  limestone  regions,  the  limestone  is 
dissolved  and  caves  are  often  formed  or  enlarged.  When 
the  water  enters  a  cave  and  drips  from  the  top,  the  water 
evaporates,  or  the  gas  escapes,  or  both,  and  the  calcium  car- 
bonate is  redeposited,  often  forming  stalactites  and  stalag- 
mites. The  stalactites  hang  from  the  roof  like  icicles,  jvvhile 


FIG.  71 .  —  Calcite  crystals. 

the  stalagmites  grow  upon  the  floor,  as  the  deposit  slowly 
accumulates  from  the  solution  which  drops  from  the  roof 
or  the  tips  of  stalactites.  The  Mammoth  Cave  in  Kentucky, 
the  Marengo  Cave  in  Indiana,  and  the  Luray  Cavern  in  Vir- 
ginia are  famous  for  these  fantastic  formations.  Mexican 
onyx  is  a  variety  of  stalagmite.  Vast  deposits  of  this  beauti- 
ful mineral  are  found  in  Algeria  and  Mexico.  It  is  translu- 
cent and  delicately  colored,  and  is  used  as  an  ornamental 
stone,  especially  for  altars,  table  tops,  mantels,  soda  foun- 
tains, and  lamp  standards.  Deposits  of  limestone  are  found 
around  many  mineral  springs.  Travertine  occurs  near  many 
springs  in  Italy.  When  fresh,  it  is  soft  and  porous1,  but  it 
soon  hardens  and  becomes  a  durable  building  stone  in  dry 
climates.  The  walls  of  many  ancient  buildings  in  Italy 
are  travertine.  Limestone  often  contains  shells  and  fossils, 
confirming  our  belief  that  limestone  is  the  remains  largely 
of  the  shells  of  animals.  The  calcium  carbonate  dissolved 


466  INORGANIC  CHEMISTRY 

in  the  ocean  is  transformed  by  marine  organisms  into  shells 
and  bony  skeletons.  The  hard  parts  of  these  animals  accu- 
mulate in  vast  quantities  on  the  ocean  bottom,  become  com- 
pact, and  are  finally  elevated  above  the  surface  of  the  ocean 
On  certain  parts  of  the  coast  of  Florida,  coquina  or  shell 
rock  is  found.  It  is  a  mass  of  fragments  of  shells  cemented 
by  calcium  carbonate,  and  in  time  it  may  become  compact 


FIG.  72.  —  Ooze  from  the  ocean  bot-       FIG.  73.  —  Globigerina  shells  (mag- 
tom,    showing    globigerina    shells  nified)  found  in  chalk  from  Iowa, 

(magnified). 

limestone.  Chalk  is  the  remains  of  shells  of  minute  animals. 
When  examined  under  the  microscope,  a  good  specimen  is 
seen  to  consist  almost  entirely  of  tiny  shells.  The  ocean 
contains  myriads  of  minute  animals,  and  when  they  die, 
their  shells,  which  are  calcium  carbonate,  sink  to  the  bottom. 
As  a  result,  the  ocean  bottom  is  partly  covered  with  a  gray 
mud,  called  globigerina  ooze.  Under  the  microscope  this 
ooze  has  the  appearance  shown  in  Figure  72,  and  when  dried 
and  compressed  it  can  hardly  be  distinguished  from  chalk. 


CALCIUM  467 

It  is  believed  that  the  immense  beds  of  chalk  found  in  Eng- 
land and  other  places  were  formed  from  this  ooze.  Some 
varieties  of  chalk  under  the  microscope  closely  resemble  the 
ooze  (Fig.  73).  Whiting  is  a  variety  of  impure  chalk;  putty 
is  a  mixture  of  whiting  and  oil.  Coral  is  calcium  carbonate, 
and  the  vast  accumulations  in  the  sea  are  the  skeletons 
of  the  coral  animals. 

The  properties  of  calcium  carbonate,  discussed  in  Chapter 
XV,  may  be  profitably  reviewed  by  the  student  at  this  point. 

Calcium  Bicarbonate  or  Acid  Calcium  Carbonate,  Ca(HCO3)2 
or  H2Ca(C03)2,  is  a  salt  formed  by  dissolving  calcium  car- 
bonate in  water  containing  carbon  dioxide.  It  decomposes 
readily  and  deposits  the  normal  carbonate  (CaCO3).  (See 
also  Stalactites  and  Hardness  of  Water.) 

Calcium  Oxide,  CaO,  is  the  chemical  name  of  lime.  It  is  a 
porous  white  solid.  Pure  lime  is  almost  infusible,  and  when 
heated  in  the  oxy hydrogen  flame  it  gives  an  intensely  bright 
light,  sometimes  called  the  "lime  light."  (See  Hydrogen  and 
Calcium  Light.)  In  the  electric  furnace  it  melts  and  vol- 
atilizes, if  the  heating  is  prolonged.  Lime  containing  im- 
purities, like  sand,  clay,  and  iron  compounds,  melts  quite 
readily  into  a  glass  or  slag.  Exposed  to  the  air,  lime  becomes 
"  air  slaked,"  that  is,  it  slowly  absorbs  water  and  carbon 
dioxide,  swells,  and  soon  crumbles  to  a  powder,  which  is  a 
mixture  of  calcium  hydroxide  and  calcium  carbonate.  When 
fresh  lime  and  water  are  mixed,  the  lime  soon  becomes 
warm,  swells,  and  finally  liberates  considerable  heat,  as  may 
often  be  seen  when  mortar  is  being  prepared.  This  opera- 
tion is  called  " slaking,"  and  the  product  is  "slaked  lime." 
The  equation  for  the  chemical  change  is  — 
CaO  +  H20  =  Ca(OH), 

Calcium  Water  Calcium 

Oxide  Hydroxide 


468 


INORGANIC  CHEMISTRY 


Fresh  lime  disintegrates  organic  matter,  and  is  therefore 
often  called  "caustic  lime"  or  quicklime.  It  combines  with 
water  to  form  calcium  hydroxide  and  interacts  with  acids 
to  form  calcium  salts. 

Lime  is  one  of  the  most  useful  substances.  It  is  used  in 
preparing  mortar,  cement,  and  metals,  in  making  bleaching 
powder,  calcium  carbide,  and  glass,  in  purifying  illuminating 
gas  and  sugar,  in  removing  hair  from  hides  before  the  process 
of  tanning,  in  dyeing  and  bleaching  cotton  cloth,  in  drying 
gases,  and  as  a  disinfectant  and  fertilizer. 

Lime  is  prepared  on  a  large  scale  by  heating  limestone  in 
a  partly  closed  cavity  or  vessel.  The  decomposition  takes 
place  according  to  the  equation  — 

CaCO3     =     CaO     +     CO, 

Calcium  Calcium  Carbon 


Carbonate 


Oxide 


Dioxide 


The  carbon  dioxide  gas  escapes,  and  the  lime  is  left  in  the 

kiln. 

Limestone  was  formerly 
"burned"  in  a  cavity  on  a 
hillside,  and  in  some  regions 
it  is  so  prepared  to-day.  This 
is  a  crude  process  and  has 
been  largely  superseded  by 
well-regulated  decomposition 
in  limekilns.  The  older  or 
periodic  kiln  is  constructed  of 
bricks  or  of  blocks  of  lime- 
stone and  loosely  filled  with 
pieces  of  limestone  from  the 

FIG.  74. -Periodic  limekiln  (vertical   arch  to  the  top>    ag  shown   jn 

Figure  74.     The  fire  is  built 

at  the  bottom  and  burns  several  days;  when  the  kiln  is  cool, 
the  lime  is  removed.    These  kilns  have  been  largely  replaced 


CALCIUM 


469 


by  continuous  kilns  (Fig.  75).  Limestone  is  introduced  at 
A  and  decomposed  by  the  heat  from  the  gases  generated  in 
B,  B.  The  lime  is  withdrawn  at  C,  C.  Carbon  dioxide 
and  other  waste  gases  escape 
through  the  top  of  the  kiln. 

Cement  is  made  by  heating 
a  pulverized,  carefully  propor- 
tioned mixture  of  limestone, 
clay,  and  sand.  This  mixture 
is  fed  into  the  upper  end  of  a 
long,  inclined,  tubular  furnace, 
which  is  heated  by  powdered 
coal  blown  in  at  the  lower 
end ;  as  the  furnace  revolves 
slowly  the  mixture  interacts, 
moves  along,  and  finally  drops 
out  as  "  clinker,"  which  con- 
sists essentially  of  calcium  sili- 
cate, calcium  aluminate,  and  an  excess  of  calcium  oxide. 
The  "  clinker  "  is  pulverized  and  mixed  with  ground  gypsum. 
Cement  and  water  slowly  interact  and  the  mixture  sets  to  a 
hard  mass ;  it  sets  under  water  as  well  as  in  air.  Concrete 
is  a  mixture  of  cement,  water,  sand,  and  crushed  stone. 
Cement  and  concrete  are  used  as  construction  material. 

Soda  Lime  is  a  mixture  of  lime  and  sodium  hydroxide. 

Calcium  Hydroxide,  Ca(OH)2,  is  a  white  powder.  It  is 
sparingly  soluble  in  water,  but  more  soluble  in  cold  than  in 
warm  water.  The  solution  has  a  bitter  taste,  an  alkaline 
reaction,  and  is  commonly  called  limewater.  Exposed  to 
the  air,  limewater  becomes  covered  with  a  thin  film  of  cal- 
cium carbonate,  owing  to  the  interaction  of  carbon  dioxide 
and  calcium  hydroxide.  For  the  same  reason,  limewater 


FIG.  75.  —  Continuous  limekiln. 


470  INORGANIC  CHEMISTRY 

becomes  turbid  or  cloudy  when  carbon  dioxide  is  passed  into 
it.  The  formation  of  calcium  carbonate  in  this  way  is  the 
usual  test  for  carbon  dioxide.  The  equation  for  this  chemical 
change  is  — 

Ca(OH)2     +     CO2     =     CaCO3     +     H20 

Calcium  Carbon  Calcium 

Hydroxide  Dioxide  Carbonate 

Limewater  is  prepared  by  carefully  adding  lime  to  consider- 
able water,  allowing  the  mixture  to  stand  for  a  day  or  two 
in  a  stoppered  bottle,  and  then  removing  the  clear  liquid. 
When  considerable  calcium  hydroxide  is  suspended  in  the 
liquid,  the  mixture  is  called  milk  of  lime.  Ordinary  white- 
wash is  thin  milk  of  lime.  Limewater  is  used  in  the  chemical 
laboratory  and  as  a  medicine. 

Mortar  is  a  thick  paste  formed  by  mixing  lime,  sand,  and 
water.  It  slowly  hardens  or  "sets,"  owing  to  the  loss  of 
water  and  the  absorption  of  carbon  dioxide.  It  hardens 
without  much  shrinking,  and  when  placed  between  bricks 
or  stones  holds  them  firmly  in  place.  The  sand  makes  the 
mass  porous  and  thus  facilitates  the  change  of  the  hydroxide 
into  the  carbonate.  The  sand  is  changed  chemically  only 
to  a  very  slight  extent,  if  at  all.  Hair  is  sometimes  added  to 
make  the  mortar  stick  better,  especially  when  it  is  used  as 
plaster  for  walls. 

Calcium  Sulphate,  CaS04.  — This  salt  occurs  abundantly 
and  in  several  varieties.  Anhydrite  (CaSO4)  is  associated 
with  sodium  chloride.  The  other  varieties  are  often  grouped 
under  the  general  term  gypsum  and  have  the  composition 
CaSO4 . 2  H2O.  Ordinary  gypsum  occurs  as  white  masses 
and  is  known  as  rock  or  massive  gypsum.  The  lustrous, 
translucent,  crystalline  kind  is  called  selenite  and  is  very 
pure;  fine-grained  masses  are  named  alabaster,  and  the 


CALCIUM  471 

fibrous  varieties  are  satin  spar.  Most  varieties  of  gypsum 
are  soft,  selenite  being  easily  scratched  with  the  finger  nail. 

Calcium  sulphate  is  slightly  soluble  in  water.  When  gyp- 
sum is  heated,  it  gradually  loses  its  water  of  crystallization, 
and  becomes  opaque  (if  previously  crystalline)  and  friable. 
When  this  dehydrated  product  is  mixed  with  water,  it  forms 
a  paste  which  soon  solidifies  or  "sets"  to  a  white  porous 
mass  with  a  smooth  surface.  If  the  raw  material  is  of  good 
quality  and  the  temperature  is  kept  near  125°  C.,  the  final 
product  is  plaster  of  Paris;  it  has  the  composition 
(CaSO4)2.H2O  and  sets  quickly  owing  to  the  formation  of  a 
network  of  crystals  of  the  less  soluble  salt  (CaSO4 .  2  H2O) . 
But  if  the  gypsum  is  impure  or  the  temperature  high,  the 
product  sets  slowly;  this  variety  is  called  cement  plaster. 
Sometimes  a  slowly  setting  plaster  is  made  by  adding  a  re- 
tarder  such  as  alum  or  borax  during  or  after  the  calcination. 
Many  kinds  of  plaster  are  made  from  gypsum  and  their  names 
are  applied  rather  indefinitely. 

Plaster  of  Paris  expands  slightly  in  setting  and  is  therefore 
used  to  make  exact  reproduction  of  statues  and  ornaments; 
considerable  is  used  in  making  molds  and  in  cementing  glass 
to  metal.  Alabaster,  being  soft  and  beautiful  in  texture, 
is  carved  into  statues  and  ornaments.  Many  grades  of  cal- 
cined gypsum  are  used  as  a  coating  for  walls  and  as  the 
principal  ingredient  of  fine  plasters.  Crude  gypsum  is  used 
in  the  manufacture  of  glass  and  porcelain;  pulverized  gypsum 
is  used  as  a  fertilizer  under  the  name  of  land  plaster.  Stucco 
is  a  mixture  of  plaster  of  Paris  and  glue. 

Calcium  Chloride,  CaCl2.  is  a  white  solid.  It  absorbs  mois- 
ture rapidly  and  is  used  to  dry  many  gases,  not  only  in  the 
laboratory,  but  also  in  such  industrial  processes  as  the  manu- 
facture of  carbon  dioxide.  The  crystallized  variety  dissolves 
readily  in  water,  and  the  solution  is  attended  by  a  marked 


472  INORGANIC  CHEMISTRY 

fall    of    temperature.     A    mixture    of    crystallized    calcium 
chloride  and  snow  produces  a  temperature  of  —  40°  C. 

The  low  freezing  point  of  calcium  chloride  solutions  has 
led  to  their  use  in  refrigerating  plants  and  in  automatic 
sprinkler  systems.  The  liquid  left  from  the  interaction  of 
calcium  carbonate  and  hydrochloric  acid  contains  calcium 
chloride,  which  on  concentration  is  deposited  in  large  crystals. 
These  readily  absorb  water,  but  lose  their  own  water  of 
crystallization  when  heated  above  200°  C.  This  anhydrous 
calcium  chloride  is  porous  and  is  the  form  usually  used  as  a 
drying  agent.  At  a  high  temperature  it  melts,  and  solidifies 
in  cooling  to  a  hard  mass  known  as  fused  calcium  chloride. 

Calcium  chloride  is  found  in  small  quantities  in  some  of  the  Stass- 
furt  salts.  It  is  obtained  in  large  quantities  as  a  by-product  in  the 
manufacture  of  sodium  carbonate  (by  the  Solvay  process)  and  other 
chemicals. 

Compounds  of  Calcium.  —  Calcium  cyanamide  (CaN2C) 
is  made  by  passing  nitrogen  over  heated  calcium  carbide ; 
it  is  used  as  a  fertilizer  because  it  provides  nitrogen  in  a  form 
readily  assimilated  by  growing  plants.  Calcium  sulphide 
(CaS)  is  formed  by  heating  a  mixture  of  gypsum  and  car- 
bon; like  other  sulphides,  it  stains  silver  brown.  Cal- 
cium oxalate  (CaC2O4)  is  a  white  solid  formed  by  the  in- 
teraction of  ammonium  oxalate  and  any  soluble  calcium 
compound ;  it  is  insoluble  in  acetic  acid  but  soluble  in 
hydrochloric  acid.  Its  formation  and  properties  serve  as 
a  test  for  calcium  compounds.  Another  test  is  the  yellow- 
red  color  imparted  to  the  Bunsen  flame  by  calcium  com- 
pounds, especially  the  chloride.  The  spectrum  of  calcium 
is  characterized  by  a  red  and  a  green  line. 

Calcium  compounds  in  aqueous  solution  yield  colorless 
calcium  ions  (Ca++).  The  atomic  weight  of  calcium  is 
40.07,  and  its  valence  is  two. 


STRONTIUM   AND   BARIUM  473 

Calcium  Compounds  and  Hardness  of  Water.  —  Calcium 
sulphate  is  slightly  soluble  in  water,  and  calcium  carbonate, 
as  we  have  already  seen,  is  changed  into  the  soluble,  unstable 
acid  carbonate  by  water  containing  carbon  dioxide.  Water 
having  these  salts  of  calcium  in  solution  is  called  hard  water. 
They  form  sticky,  insoluble  compounds  with  soap,  and  as  long 
as  water  contains  such  salts,  the  soap  is  useless  as  a  cleansing 
agent.  Heat  decomposes  acid  calcium  carbonate,  and  the 
hardness  due  to  the  acid  carbonate  is  called  temporary  hard- 
ness, because  boiling  removes  it;  temporary  hardness  can 
also  be  removed  by  the  addition  of  some  slaked  lime.  The 
hardness  caused  by  calcium  sulphate  cannot  be  removed 
by  boiling  and  is  called  permanent  hardness.  Magnesium 
sulphate,  like  calcium  sulphate,  also  produces  permanent 
hardness.  Permanent  hardness  can  be  removed,  however, 
by  adding  sodium  carbonate  to  the  water,  because  the  cal- 
cium sulphate  and  sodium  carbonate  interact  and  produce 
soluble  sodium  sulphate  and  insoluble  calcium  carbonate; 
the  latter  can  be  removed  by  filtering.  On  a  small  scale 
borax  or  ammonia  may  be  used  to  remove  both  kinds  of 
hardness.  When  hard  water  is  used  in  steam  boilers,  the 
calcium  salts  are  often  deposited  as  a  hard  mass  known  as 
"  boiler  scale."  Soft  water,  such  as  rain  water,  contains 
little  or  no  calcium  or  magnesium  salts. 

STRONTIUM  AND  BARIUM 

Strontium,  Sr,  and  Barium,  Ba,  are  uncommon  metallic 
elements.  They  resemble  calcium  closely  in  their  physical 
and  chemical  properties.  The  metals  themselves  never 
occur  free,  and  are  hardly  more  than  chemical  curiosities. 
Their  compounds  are  abundant,  and  some  are  useful. 

Compounds  of  Strontium.  —  The  native  compounds  are 
the  beautifully  crystalline  minerals,  strontianite  (strontium 


474  INORGANIC  CHEMISTRY 

carbonate,  SrCO8)  and  celestite  (strontium  sulphate,  SrS04). 
Strontium  oxide  (strontia,  SrO),  like  lime,  is  made  by  heating 
the  carbonate.  It  unites  with  water  to  form  strontium 
hydroxide  (Sr(OH)2),  which  is  used  in  the  purification  of 
beet  sugar.  Strontium  nitrate  (Sr(NO3)2)  and  other  salts 
of  strontium  color  a  flame  crimson,  and  are  widely  used  in 
making  fireworks,  especially  "red  fire."  The  latter  is  a 
mixture  of  potassium  chlorate,  shellac,  and  strontium  nitrate. 
Several  strontium  salts  are  used  as  medicines. 

The  crimson  color  imparted  to  the  Bunsen  flame  is  the 
usual  test  for  strontium.  The  spectrum  of  strontium  is  char- 
acterized by  several  red  lines  and  a  blue  one. 

Compounds  of  Barium.  —  The  most  abundant  native 
compounds  are  witherite  (barium  carbonate,  BaCO3)  and 
barite  (barium  sulphate,  heavy  spar,  barytes,  BaSO4).  The 
oxides,  BaO  and  BaO2,  have  already  been  mentioned  as  a 
source  of  oxygen.  Barium  hydroxide  (Ba(OH)2)  solution  is 
often  called  baryta  water,  and  it  forms  insoluble  barium  car- 
bonate (BaCO3)  by  interaction  with  carbon  dioxide.  Barium 
chloride  (BaCl2)  is  a  white  crystalline  solid.  Barium  sul- 
phate (BaSO4)  is  formed  as  a  fine,  white,  highly  insoluble 
precipitate  by  the  combination  of  barium  and  sulphate  ions. 
It  is  used  in  making  (and  sometimes  adulterating)  white 
paint.  A  specially  prepared  mixture  of  barium  sulphate 
and  zinc  sulphide,  called  lithophone,  forms  the  basis  of  a 
white  paint  which  is  superior  in  some  ways  to  white  lead 
paint.  Barium  sulphate  is  also  used  to  fill  paper  and  give 
it  a  gloss.  Barium  salts  color  a  flame  green,  and  barium 
nitrate  (Ba(NO3)2)  is  extensively  used  in  making  fireworks, 
especially  "  green  fire."  Commercial  barium  sulphide  (BaS), 
as  well  as  the  sulphides  of  calcium  and  strontium,  shine  feebly 
in  the  dark,  after  having  been  exposed  to  a  bright  light.  On 
account  of  this  property  they  are  used  in  making  luminous 


STRONTIUM  AND  BARIUM  475 

paint.  Barium  chromate  (BaCrO4)  is  a  yellow  solid  ob- 
tained by  the  interaction  of  a  soluble  barium  compound  and 
potassium  dichromate.  Soluble  barium  salts,  unlike  those 
of  calcium,  are  poisonous. 

The  green  color  imparted  to  the  Bunsen  flame  is  the  test 
for  barium.  The  spectrum  of  barium  is  characterized  by 
several  green  and  orange  lines. 

Aqueous  solutions  of  strontium  and  of  barium  compounds 
contain  colorless  strontium  ions  (Sr++)  and  barium  ions 


The  Alkaline  Earth  Family.  —  This  is  a  typical  family 
and  resembles  its  contiguous  families,  the  alkali  metals  and 
the  earth  metals.  The  metals  themselves,  if  calcium  is  taken 
as  a  type,  are  less  active  than  sodium  and  potassium  but 
more  active  than  aluminium.  Not  only  have  these  elements 
properties  which  are  much  alike,  but  they  also  show  a  grada- 
tion in  properties  as  the  atomic  weight  increases.  All  interact 
with  cold  water,  burn  in  air,  and  form  analogous  compounds 
whose  properties  are  strikingly  suggestive  of  family  relation- 
ship; on  the  other  hand  their  densities  increase  from  about 
1.5  for  calcium  (at.  wt.  40.07)  through  2.5  for  strontium  (at. 
wt.  87.63)  to  3.57  fcr  barium  (at.  wt.  137.37).  All  three  ele- 
ments form  an  hydroxide,  a  carbonate,  a  nitrate,  and  a  sul- 
phate. The  solubility  of  the  chlorides  and  nitrates  is  quite 
marked  and  decreases  in  the  order  in  which  the  metals  have 
been  studied  (calcium,  strontium,  barium);  the  hydroxides 
are  much  less  soluble  and  in  the  opposite  order;  while  the 
sulphates  vary  widely  in  solubility,  a  liter  of  water  dissolving 
about  2  gm.  of  calcium  sulphate  and  only  .0023  gm.  of  barium 
sulphate.  All  the  compounds  of  these  elements  are  white, 
except  barium  chromate,  which  is  yellow.  The  valence  of 
the  alkaline  earth  elements  in  their  compounds  is  almost 
invariably  two. 


476  INORGANIC  CHEMISTRY 

RADIUM 

Radium  is  a  rare  element.  It  is  a  constituent  of  certain 
rare  uranium-bearing  minerals,  especially  pitchblende  and 
carnotite.  Pitchblende  is  found  in  Bohemia  and  carnotite 
in  Colorado  and  Utah. 

The  proportion  of  radium  in  pitchblende  and  carnotite  is 
minute,  only  a  few  milligrams  to  the  ton.  Nevertheless  the 
radium  is  extracted  from  these  minerals  and  crystallized  as 
radium  bromide.  The  supply  of  radium  ores  is  very 
limited. 

The  general  properties  of  radium  compounds  show  that 
the  element  belongs  to  the  alkaline  earth  group.  Metallic 
radium,  which  was  first  isolated  by  Madame  Curie  in  1910, 
closely  resembles  barium.  It  is  a  silvery  white,  intensely 
active  metal.  Radium  forms  compounds  analogous  to  those 
of  barium,  the  best  known  being  the  chloride  (RaCl2),  bro- 
mide (RaBr2),  and  the  sulphate  (RaSO4).  The  bromide  is 
the  common  commercial  salt  and  is  the  substance  usually 
meant  by  the  term  radium.  Radium  compounds  color  the 
Bunsen  flame  red,  have  a  different  spectrum  from  other  ele- 
ments, are  naturally  phosphorescent,  and  produce  phos- 
phorescence in  various  substances ;  e.g.  barium  platino- 
cyanide  (BaPt(CN)4),  diamond,  willemite,  and  zinc  sulphide. 
Besides  the  properties  just  mentioned,  radium  compounds 
have  others  which  are  characteristic  and  differ  from  those 
exhibited  by  most  substances.  Thus,  they  spontaneously 
evolve  relatively  large  quantities  of  heat,  affect  a  photo- 
graphic plate,  and  discharge  an  electroscope.  These  proper- 
ties are  called  radioactive  properties.  Similar,  though  less 
active,  properties  are  possessed  by  uranium,  thorium,  and 
other  substances;  and  radioactivity  can  be  more  appro- 
priately discussed  under  these  elements.  (See  Chapter 
XXXII.) 


RADIUM  477 

PROBLEMS  AND  EXERCISES 

1.  If  the  specific  gravity  of  marble  is  2.75,  how  many  grams 
would  a  cubic  meter  weigh? 

2.  What  weight  of  barium  sulphate  can  be  obtained  by  the 
interaction  of  barium  chloride  (BaCl2)  and  10  gm.  of  crystallized 
magnesium  sulphate,  MgSO4 .  7  H2O  ? 

3.  An  analysis  of  a  sample  of  limestone  gave  96.45  per  cent 
CaCO3,  1.00  per  cent  Si02,  .78  per  cent  MgO,  1.76  per  cent  Fe2O3 
and  A1203.     How  much  pure  lime  could  be  made  from  100  tons  of 
this  limestone? 

4.  What  is  (a)  the  weight  and  (6)  the  volume  (at  standard  con- 
ditions) of  the  gas  liberated  by  the  interaction  of  water  and  10  gm. 
of  calcium? 

5.  Write  the  formulas  of  the  following  compounds  by  applying 
the  principle  of  valence  (Chapter  XIV) :   Acid  calcium  phosphate, 
calcium  nitrate,  calcium  iodide,  calcium  bromide,  strontium  chloride, 
strontium  oxide,  barium  sulphide,  barium  monoxide,  barium  chlo- 
rate, calcium  phosphate  (ortho),  strontium  sulphide. 

6.  Apply  Exercise  5  to  radium. 

7.  Calculate  the  atomic  weight  of  radium  if  2.00988  mg.  of 
radium  chloride  are  formed  by  2.61099  mg.  of  radium  bromide. 

8.  Calculate  the  solubility  product  of  calcium  oxalate  if  the 
ionization  is  96  per  cent  and  the  molar  solubility  is  .000044. 

9.  As  in  8,  of  calcium  sulphate  for  52.5  per  cent  and  .015. 


CHAPTER  XXVIII 
Magnesium,  Zinc,  Cadmium,  and  Mercury 

THESE  elements  form  a  natural  family  in  Group  II  of  the 
periodic  classification,  though  the  members  are  not  as 
closely  related  as  in  the  alkali  and  alkaline  earth  families. 

MAGNESIUM 

Occurrence  of  Magnesium.  —  Magnesium  is  never  found 
free.  In  combination  it  is  widely  distributed  and  very 
abundant,  constituting  about  2.5  per  cent  of  the  earth's 
crust.  Dolomite  is  magnesium  calcium  carbonate 
(CaCO8 .  MgCO8  or  CaMg(CO3)2);  it  forms  whole  mountain 
ranges  in  the  Tyrol  and  vast  deposits  in  many  regions. 
Dolomite  resembles  marble  and  limestone  in  its  properties. 
Magnesium  carbonate  (MgCO8)  is  also  abundant.  Many  of 
the  Stassfurt  salts  contain  magnesium,  for  example,  kainite 
(KC1,  MgSO4 .  3  H2O),  carnallite  (KC1,  MgCl, .  6  H2O),  and  kie- 
serite  (MgSO4.  H2O).  Magnesium  is  a  component  of  serpen- 
tine, talc,  soapstone,  asbestos,  meerschaum,  and  other  sili- 
cates. The  sulphate  (MgSO4)  and  chloride  (MgCl2)  are  found 
in  sea  water  and  in  mineral  springs. 

Preparation  of  Magnesium.  —  Magnesium  was  formerly 
prepared  by  reducing  the  chloride  with  sodium.  It  is  now 
economically  manufactured  by  the  electrolysis  of  fused 
carnallite.  A  sketch  of  the  essential  parts  of  the  apparatus 
is  shown  in  Figure  76.  Carnallite  is  put  into  the  cylindrical 
iron  vessel  C,  which  is  the  cathode.  This  is  closed  by  the 

478 


MAGNESIUM 


479 


E 


j^di.MJ^    ' 


B 


air-tight  cover  through  which  pass    the    pipes  D,  D'    for 

conveying  inert  gases  into  and  out  of  the  apparatus.     The 

carbon   anode    A   dips  into  the 

carnallite     and    is    inclosed    by 

the  porcelain  cylinder  B,  which 

is  provided  with  a  pipe  E,  for  the 

escape  of  the  chlorine  liberated 

at  the  anode.     The  carnallite  is 

kept    fused    by    external    heat. 

When   the    current    passes,    the 

chlorine  liberated  at  the  anode 

escapes  through  E,  and  the  mag- 

nesium   liberated   at   the    Cathode 
a  ,1       £         i  I-,.,  j 

floats  on  the  fused  carnallite  and 


FIG.  76.  —  Apparatus  for  the 

manufacture  of  magnesium  by 
the  electrolysis  of  carnallite. 


is  prevented  from  oxidizing  by  the  inert  gas  supplied  through 
D.  The  porcelain  cylinder  B  prevents  the  chlorine  from 
escaping  into  the  larger  vessel.  The  molten  magnesium  is 
carefully  removed  at  intervals. 

Properties  of  Magnesium.  —  Magnesium  is  a  lustrous,  sil- 
very white  metal.  It  is  a  light  metal,  the  specific  gravity  be- 
ing only  about  1.75.  It  is  tenacious  and  ductile,  and  when 
hot  can  be  drawn  into  wire  or  pressed  into  ribbon,  the  latter 
being  a  common  commercial  form.  It  melts  at  about  650°  C. 
and  can  be  cast  into  different  shapes.  At  a  higher  tempera- 
ture it  volatilizes.  It  burns  with  a  dazzling  light,  producing 
dense  white  clouds  of  magnesium  oxide  (MgO)  together  with 
some  magnesium  nitride  (Mg3N2).  It  does  not  tarnish  in 
dry  air,  but  in  moist  air  it  is  soon  covered  with  a  film  of 
oxide.  It  liberates  hydrogen  from  acids  and  from  boiling 
water.  Heated  in  nitrogen,  it  forms  magnesium  nitride 
(Mg3N2). 

Uses  of  Magnesium.  —  Magnesium  in  the  form  of  powder 
is  used  chiefly  in  taking  flash-light  photographs.  Small 


480  INORGANIC  CHEMISTRY 

quantities  are  used  in  making  fireworks.  The  powder  and 
ribbon  are  used  in  the  chemical  laboratory  as  a  convenient 
form  of  the  metal.  Some  is  used  to  reduce  rare  metals 
from  their  oxides. 

Magnesium  Oxide,  MgO,  is  a  white,  bulky  powder.  It  is 
formed  when  magnesium  burns  in  the  air,  but  it  is  manu- 
factured by  gently  heating  magnesite  (MgC03),  just  as  lime 
is  made  from  limestone.  It  is  often  called  magnesia,  or 
calcined  magnesia.  The  native  oxide  is  the  mineral  periclase. 
Magnesia  dissolves  very  slightly  in  water,  forming  magnesium 
hydroxide  (Mg(OH)2).  A  mixture  of  magnesia  and  water, 
with  or  without  magnesium  chloride,  hardens  on  exposure 
to  the  air,  and  is  sometimes  used  as  a  cement  or  artificial 
stone.  Native  magnesium  hydroxide  is  the  mineral  brucite. 
Like  lime,  magnesia  withstands  a  high  temperature,  and  is, 
therefore,  used  as  the  lining  of  electrothermal  apparatus, 
metallurgical  furnaces,  and  cement  kilns,  and  as  the  chief 
ingredient  of  a  protective  covering  for  steam  pipes.  Mag- 
nesia is  used  in  medicine  as  a  mild  alkali  and  as  an  antidote 
for  poisoning  by  mineral  acids. 

Magnesium  Sulphate,  MgSO4,  is  a  white  solid.  There  are 
several  crystalline  varieties.  The  native  salt  kieserite 
(MgSO4.H2O),  when  added  to  water,  changes  into  Epsom 
salts  (MgSO4.7H2O).  This  variety  was  first  found  in  the 
mineral  spring  at  Epsom,  England.  It  is  efflorescent,  very 
soluble  in  water,  and  its  solution  has  a  bitter  taste.  It  is 
used  in  medicine  as  a  purgative,  in  manufacturing  sulphates 
of  sodium  and  potassium,  as  a  fertilizer  in  place  of  gypsum, 
and  as  a  coating  for  cotton  cloth. 

Magnesium  Chloride,  MgCl2,  is  a  white  solid.  It  is  a  by- 
product in  the  preparation  of  potassium  chloride.  The 
crystallized  salt  (MgCl2 .  6  H2O)  is  very  deliquescent.  Mag- 
nesium chloride  hydrolyzes  with  hot  water,  forming  magne- 


MAGNESIUM  481 

slum  hydroxide  and  hydrochloric  acid.  Hence  water  con- 
taining magnesium  chloride  (e.g.  sea  water)  is  not  suitable 
for  use  in  boilers. 

Magnesium  Carbonate,  MgCO3,  occurs  native  as  magnesite, 
and  combined  with  calcium  carbonate  as  dolomite.  Mag- 
nesite is  converted  by  heat  into  magnesia  and  carbon  diox- 
ide ;  the  magnesia,  as  stated  above,  is  very  generally  used  as 
a  refractory  material,  while  the  gas  is  utilized  in  the  prepara- 
tion of  liquid  carbon  dioxide  or  of  charged  beverages.  The 
commercial  salt  known  as  magnesia  alba,  or  simply  mag- 
nesia, is  a  basic  carbonate  (Mg4(OH)2(C03)3 . 3  H20). 

It  was  during  an  investigation  of  magnesia  alba  that  Black  (1728- 
1799)  discovered  carbon  dioxide  and  showed  the  close  relation  between 
analogous  compounds  of  magnesium  and  calcium. 

Miscellaneous.  —  Besides  the  oxide  and  sulphate  other 
magnesium  compounds  are  used  as  medicines.  Fluid  mag- 
nesia is  prepared  by  dissolving  magnesium  carbonate  in 
water  containing  carbon  dioxide  or  by  suspending  mag- 
nesium hydroxide  in  water;  it  is  a  mild  alkali  and  laxative. 
Magnesium  citrate  has  a  similar  action;  it  is  an  effervescing 
mixture  prepared  from  sodium  bicarbonate,  tartaric  and 
citric  acids,  sugar,  and  magnesium  sulphate.  When  disodium 
phosphate  and  ammonium  hydroxide  are  added  to  the  solu- 
tion of  a  magnesium  compound,  a  white  crystalline  precipi- 
tate of  ammonium  magnesium  phosphate  (NH4MgP04)  is 
produced;  its  formation  is  a  test  for  magnesium.  Another 
test  (though  often  indifferent)  is  the  pale  pink  residue  left 
after  a  magnesium  compound  has  been  intensely  heated  in 
a  blowpipe  flame  and  subsequently  moistened  with  cobalt- 
ous  nitrate  solution. 

Soluble  magnesium  compounds  yield  colorless  magnesium 
ions  (Mg++).  The  atomic  weight  of  magnesium  is  24.32, 
and  the  valence  is  two. 


482  INORGANIC   CHEMISTRY 

ZINC 

Occurrence  of  Zinc.  — Free  zinc  is  never  found.  The  ores 
of  zinc  are  not  numerous,  though  rather  abundant.  The 
chief  gres  are  the  sulphide  (sphalerite  or  zinc  blende,  ZnS), 
the  carbonate  (smithsonite,  ZnCO3),  the  silicate  (calamine, 
H2Zn2SiO5),  and  the  red  oxide  (zincite,  ZnO).  Franklinite 
(Zn(FeO2)2)  and  willemite  (Zn2SiO4)  are  complex  ores  which 
also  contain  manganese  and  iron.  Gahnite  has  the  com- 
position ZnAl2O4. 

Zinc  ores  are  found  in  Germany,  Italy,  France,  Greece,  Austria- 
Hungary,  Belgium,  England,  and  the  United  States.  Missouri  and 
Kansas  contain  large  deposits  of  the  sulphide,  while  the  other  ores 
occur  chiefly  in  New  Jersey. 

Metallurgy  of  Zinc.  —  Zinc  is  easily  smelted.  The  ores 
are  first  roasted  to  change  them  into  the  oxide,  thus :  — 

ZnCO3     =     ZnO      +      CO2 

Zinc  Zinc  Carbon 

Carbonate  Oxide  Dioxide 

2  ZnS     +     3O2     =     2  ZnO     +     2SO2 

Zinc  Oxygen  Zinc  Sulphur 

Sulphide  Oxide  Dioxide 

The  oxide  is  then  reduced  by  heating  it  with  coal.  This 
operation  is  conducted  in  fire-clay  retorts  connected  with 
iron  or  clay  receivers  into  which  the  vapor  passes;  at  first 
it  condenses  as  a  blue-gray  powder  known  as  zinc  dust, 
which  is  really  a  mixture  of  zinc  and  zinc  oxide;  but  it  finally 
condenses  as  a  liquid,  which  is  drawn  off  at  intervals  and  cast 
into  bars  or  plates.  The  impure,  massive  zinc  thus  obtained 
is  called  spelter;  it  is  freed  from  carbon,  lead,  iron,  cadmium, 
and  arsenic  by  repeated  distillation,  often  under  reduced 
pressure. 

Properties  of  Zinc.  —  Zinc  is  a  bluish  white,  lustrous  metal. 
Its  physical  properties  vary  with  the  temperature.  At  or- 


ZINC  483 

dinary  temperatures  it  is  brittle,  but  at  120°-150°  C.  it  is 
soft  and  can  be  rolled  into  sheets  and  drawn  into  wire,  while 
its  specific  gravity  rises  from  6.9  to  7.2.  Zinc  which  has 
been  rolled  or  drawn  does  not  become  brittle  upon  cooling, 
but  remains  pliable.  At  200°  C.  it  again  becomes  brittle 
and  can  be  easily  pulverized.  It  melts  at  about  420°  C. 
and  boils  at  about  920°  C.  Heated  in  the  air  above  its 
melting  point,  zinc  burns  with  a  bluish  green  flame,  forming 
white  zinc  oxide  (ZnO).  Zinc  does  not  tarnish  in  dry  air, 
but  in  moist  air  it  slowly  becomes  coated  with  a  hard,  co- 
herent film  which  prevents  further  action.  With  hot  solu- 
tions of  sodium  and  potassium  hydroxides,  it  forms  zincates 
and  liberates  hydrogen,  thus :  — 

2KOH     +     Zn     =     H2     +     K2Zn02 

Potassium  Zinc          Hydrogen        Potassium 

Hydroxide  Zincate 

Commercial  zinc  interacts  with  dilute  acids  and  liberates 
hydrogen  (except  from  nitric  acid).  Pure  zinc  interacts 
with  acids  if  in  contact  with  platinum,  iron,  or  copper,  or  if 
copper  sulphate  is  added  to  the  mixture.  Zinc  displaces 
many  other  metals  (e.g.  lead,  tin,  copper,  and  mercury)  from 
their  solutions;  it  is  strongly  electropositive.  (See  Displace- 
ment of  Metals,  Chapter  XXVI.) 

The  vapor  density  of  zinc  requires  the  molecular  weight 
66  (approximately).  Since  the  atomic  weight  is  65.37,  a 
molecule  of  zinc  vapor  contains  only  one  atom. 

Uses  of  Zinc.  — Zinc  in  the  form  of  sticks  and  plates  is 
extensively  used  as  the  positive  electrode  in  electric  batteries. 
Sheet  zinc  is  used  as  a  lining  for  tanks,  and  as  a  protective 
covering  which  is  often  placed  behind  and  beneath  stoves. 
Iron  dipped  into  zinc  becomes  coated  with  zinc  and  is  called 
galvanized  iron;  it  does  not  rust  easily,  and  is  often  used  in 
place  of  zinc  for  roofs,  pipes,  cornices,  and  water  tanks. 


484  INORGANIC  CHEMISTRY 

Telegraph  wire  is  also  galvanized.  Zinc  dust  is  sometimes 
used  in  the  .cyanide  process  of  extracting  gold  and  in  many 
chemical  experiments  in  the  laboratory.  Brass,  German 
silver,  and  other  alloys  contain  zinc.  (See  Alloys  of  Copper.) 
Antifriction  metals,  which  are  used  for  bearings,  are  alloys 
of  zinc;  Babbitt's  metal,  for  example,  contains  approxi- 
mately 69  per  cent  of  zinc,  19  of  tin,  4  of  copper,  3  of  anti- 
mony, and  5  of  lead. 

Compounds  of  Zinc.  —  Native  zinc  oxide  (ZnO)  is  red, 
owing  to  the  presence  of  manganese,  but  the  pure  oxide 
is  white  when  cold  and  yellow  when  hot.  It  is  formed  when 
zinc  burns,  and  is  manufactured  in  this  way  or  by  heating 
zinc  carbonate.  It  is  often  called  "zinc  white"  or  " Chinese 
white,"  and  is  used  to  make  a  white  paint  which  is  not  dis- 
colored by  sulphur  compounds  (especially  hydrogen  sulphide) , 
and  is  therefore  well  suited  for  painting  the  walls  of  a  chemi- 
cal laboratory.  It  is  also  used  as  an  ingredient  of  ointments 
and  cosmetics.  Native  zinc  sulphide  (ZnS)  is  yellow,  brown, 
or  black  on  account  of  impurities,  but  the  pure  sulphide  is 
white.  The  latter  is  formed  as  a  jellylike  precipitate  when 
hydrogen  sulphide  is  passed  into  an  alkaline  solution  of  a 
zinc  salt ;  it  dissolves  in  hydrochloric  acid,  but  not  in  acetic. 
Zinc  sulphide  is  used  in  making  paint  (see  page  474).  Zinc 
sulphate  (ZnSO4)  is  formed  by  the  interaction  of  zinc  and 
dilute  -sulphuric  acid.  Large  quantities  are  made  by  roast- 
ing the  sulphide  in  a  limited  supply  of  oxygen  and  extract- 
ing the  sulphate  with  water.  It  is  a  white,  crystalline  solid 
(ZnSO4.  7  H2O),  which  effloresces  in  the  air,  and  when  heated 
to  100°  C.  loses  most  of  its  water  of  crystallization.  The 
crystallized  salt  is  called  white  vitriol.  It  is  used  in  dyeing 
and  calico  printing,  as  a  disinfectant,  and  as  a  medicine.  It 
is  poisonous,  but  can  be  safely  used  externally  to  relieve  in- 
flammation. Zinc  chloride  (ZnCl2)  is  a  white,  deliquescent 


ZINC  485 

solid,  prepared  by  dissolving  zinc  in  hydrochloric  acid  and 
evaporating  the  solution  until  a  sample  solidifies  on  cooling. 
It  is  used  in  surgery,  and  also  as  a  constituent  of  a  mixture 
for  filling  teeth ;  large  quantities  are  used  to  preserve  wood, 
especially  railroad  ties.  By  hydrolysis  it  forms  hydrochloric 
acid  and  basic  zinc  chloride  (Zn(OH)Cl).  Zinc  hydroxide 
(Zn(OH)2)  is  formed  as  a  dull  white,  flocculent  precipitate 
by  the  interaction  of  sodium  or  potassium  hydroxide  and  the 
solution  of  a  zinc  salt,  thus :  — 

ZnSO4     +     2NaOH     =     Zn(OH)2     +     Na2SO4 

Zinc  Sodium  Zinc  Sodium 

Sulphate  Hydroxide  Hydroxide  Sulphate 

An  excess  of  the  alkaline  hydroxide  changes  the  zinc  hydrox- 
ide into  a  soluble  zincate,  thus :  — 

Zn(OH)2     +     2NaOH     =     Na2Zn02     +     2H2O 

Zinc  Sodium  Sodium  Water 

Hydroxide  Hydroxide  Zincate 

Zinc  salts  are  poisonous. 

Tests  for  Zinc.  —  The  formation  of  the  sulphide  or  the 
hydroxide  and  a  soluble  zincate  as  above  described  serves 
as  the  test  for  zinc.  A  green  incrustation  is  produced  when 
zinc  compounds  are  heated  on  charcoal  and  then  moistened 
with  a  cobaltous  nitrate  (Co(NO3)2)  solution. 

Miscellaneous.  —  Zinc  acts  as  a  metal  and  a  non-metal. 
Thus,  in  compounds  like  zinc  sulphate,  zinc  chloride,  and 
zinc  sulphide  it  is  a  metal;  solutions  of  the  class  contain 
colorless  zinc  ions  (Zn++).  In  the  zincates,  such  as  sodium 
zincate  (Na2ZnO2),  the  zinc  acts  as  a  non-metal.  Zinc  forms 
complex  compounds  with  ammonium  hydroxide,  e.g.  ammo- 
nio-zinc  hydroxide  (Zn(NH3)4(OH)2).  In  solutions  of  com- 
plex zinc  compounds,  the  zinc  is  often  a  part  of  a  complex  ion. 

The  valence  of  zinc  is  two. 


INORGANIC  CHEMISTRY 


CADMIUM 

Cadmium,  Cd,  is  an  uncommon  metal;  certain  compounds 
are  frequently  found  in  zinc  and  other  ores.  It  also  occurs 
as  a  sulphide  (greenockite,  CdS).  It  is  white,  lustrous,  and 
rather  soft.  Its  specific  gravity  is  about  8.6  and  its  melting 
point  is  about  320°  C.  Cadmium  is  a  constituent  of  certain 
dental  amalgams  and  fusible  alloys.  (See  Bismuth.)  Wood's 
metal  contains  12  per  cent  of  cadmium. 

The  most  important  compound  is  cadmium  sulphide  (CdS). 
This  is  a  bright  yellow  solid,  formed  by  adding  hydrogen 
sulphide  to  the  solution  of  a  cadmium  compound.  It  is 
used  as  a  pigment.  Its  formation  also  serves  as  a  test  for 
cadmium.  Warm  dilute  sulphuric  acid  dissolves  cadmium 
sulphide  owing  to  the  formation  of  soluble  cadmium  sulphate, 
but  cadmium  sulphide  does  not  dissolve  in  potassium  cyanide 
solution;  both  operations  permit  the  more  or  less  complete 
separation  of  cadmium  and  copper  compounds.  Cadmium 
hydroxide  (Cd(OH)2)  is  a  white  solid  formed  upon  the  addi- 
tion of  sodium  hydroxide  to  the  solution  of  a  cadmium  com- 
pound; it  is  insoluble  in  an  excess  of  sodium  hydroxide  but 
with  ammonium  hydroxide  it  forms  soluble  Cd(NH3)4(OH)2. 

Cadmium  has  a  valence  of  two  in  all  its  compounds.  Its 
vapor  density  requires  the  molecular  weight  112  (approxi- 
mately). Since  the  atomic  weight  is  112.4,  a  molecule  of 
the  vapor  contains  one  atom. 

Aqueous  solutions  of  cadmium  compounds  contain  color- 
less cadmium  ions  (Cd++).  The  element  also  forms  com- 
plex ions,  e.g.  Cd(NH3)4++  and  Cd(CN)4--. 

MERCURY 

Occurrence  of  Mercury.  —  Native  mercury  is  occasionally 
found  in  minute  globules,  but  the  most  abundant  ore  is 
mercuric  sulphide  (cinnabar,  HgS).  The  ore  is  mined  in 


MERCURY  487 

Spain,  Austria,  Russia,  Italy,  and  Mexico;  in  the  United 
States  large  quantities  are  obtained  in  California  and  Texas. 

Mercury  has  been  known  for  ages  as  quicksilver.  The  Latin  name, 
hydrargyrum,  which  gives  the  symbol  Hg.  means  literally,  "water 
silver,"  emphasizing  the  fact,  so  well  known,  that  mercury  looks  like 
silver  and  flows  like  water. 

Preparation  of  Mercury.  —  Mercury  is  prepared  by  roasting 
cinnabar  in  an  open  furnace  or  closed  retort.  In  the  fur- 
nace the  sulphide  is  transformed  by  the  oxygen  of  the  air 
into  mercury  and  sulphur  dioxide,  thus :  — 

HgS     +     02     =     Hg     +     S02 

Cinnabar         Oxygen       Mercury         Sulphur 

Dioxide 

In  the  retort  lime  is  mixed  with  the  ore  and  the  equation  for 
the  reaction  is  :  — 

4  HgS     +     4CaO     =     4Hg     +     3CaS     +     CaSO4 

Mercuric  Lime  Mercury  Calcium  Calcium 

Sulphide  Sulphide  Sulphate 

In  each  process  the  mercury  is  liberated  as  a  vapor  and  con- 
densed in  a  series  of  cooled  chambers.  The  crude  mercury 
is  freed  from  soot  and  dust  and  collected  into  large  globules 
by  stirring  and  rubbing;  it  is  farther  purified  by  filtering  it 
through  charcoal  or  chamois  skin.  Metallic  impurities  are 
removed  by  distillation  or  by  •  agitation  with  dilute  nitric 
acid.  Mercury  is  sent  into  commerce  in  strong  iron  flasks, 
holding  about  75  pounds. 

Properties  of  Mercury.  —  Mercury  is  a  bright,  silvery  metal, 
and  is  the  only  one  which  is  liquid  at  ordinary  temperatures. 
It  solidifies  at  about  —  38.7°  C.,  and  boils  at  about  357°  C. 
It  is  a  heavy  metal,  the  specific  gravity  being  about  13.59. 
It  is  slightly  volatile  even  at  ordinary  temperatures,  and  the 
vapor  is  poisonous.  Mercury  does  not  tarnish  in  the  air, 


488  INORGANIC   CIIKMISTHY 

unless  sulphur  compounds  are  present,.  At  a  high  tempera- 
ture it  combines  slowly  with  oxygen  to  form  the  red  oxide 
(HgO).  Hydrochloric  acid  and  cold  sulphuric  acid  do  not 
affect  it;  hot  concentrated  sulphuric  acid  oxidizes  it,  and 
nitric  acid  changes  it  into  nitrates. 

The  vapor  density  of  mercury  requires  the  molecular  weight 
200  (approximately),  Since  the  atomic  weight  is  200,  a 
molecule  of  the  vapor  contains  only  one  atom. 

Uses  of  Mercury.  — Mercury  is  used  in  making  thermom- 
eters, barometers,  manometers,  and  many  other  kinds  of 
scientific  apparatus.  Its  extensive  use  in  extracting  gold  and 
silver  has  been  mentioned.  (See  Amalgamation.)  Consider- 
able is  used  in  preparing  amalgams,  medicines,  and  explo- 
sives (e.g.  fulminating  mercury,  which  is  used  in  percussion 
caps  and  cartridges).  It  has  also  come  into  use  recently  as 
an  electrode  in  various  electrochemical  processes. 

Amalgams  are  alloys  of  mercury  with  other  metals.  They 
are  easily  prepared  by  mixing  the  constituents.  Some- 
times the  union  is  violent,  as  in  the  preparation  of  sodium 
amalgam.  Amalgamated  zinc  is  usually  used  in  electric 
batteries  to  prevent  unnecessary  loss  of  the  zinc.  Tin  amal- 
gam is  sometimes  used  to  coat  mirrors.  Amalgams  of  cer- 
tain metals  are  used  as  a  filling  for  teeth.  Care  should  be 
taken,  while  handling  mercury,  not  to  let  it  come  in  contact 
with  gold  rings,  since  gold  amalgam  is  readily  formed.  Iron 
is  the  only  common  metal  which  does  not  form  an  amalgam. 

Compounds  of  Mercury.  —  Mercury,  like  copper,  forms  two 
classes  of  compounds  —  the  mercurous  and  the  mercuric. 
The  valence  of  mercury  is  one  in  the  mercurous  compounds 
and  two  in  the  mercuric.  Mercuric  oxide  (HgO)  is  a  red 
powder,  produced  l>\  heating  mercury  in  air  or  by  heating 
a  mixture  of  mercury  and  mercuric  nitrate.  As  we  have. 


MKIM'UL'Y  -1S9 

:il  ready  seen,  mercuric  oxide  is  decomposed  by  heat  into 
mercury  and  oxygen.  A  yellow  variety  is  produced  by  the 
interaction  of  sodium  hydroxide  and  a.  meiv.urie  salt,  thus:  — • 

2  NaOH  +  Hg(N08)2   =    HgO    +    2NaN08    +    HaO 

Sodium  Mercuric  Mercuric  Sodium  Water 

Hydroxido  Nitmto  Oxide  Nitrate 

Mercurous  chloride  (Hg2Cla  or  HgCl)  is  a  white,  tasteless 
powder,  insoluble  in  water.  It  is  formed  when  a  chloride 
and  mercurous  nitrate  interact,  but  it  is  manufactured  usu- 
ally by  heating  a  mixture  of  mercuric  chloride  and  mercury. 
Under  the  name  of  calomel  it  is  extensively  used  as  a  medi- 
cine. Mercuric  chloride  (HgCla)  is  a  white,  crystalline  solid, 
soluble  in  water  and  in  alcohol.  It  is  prepared  by  heating  a 
mixture  of  mercuric  sulphate  and  common  salt.  It  is  a 
violent  poison.  The  best  antidote  is  the  white  of  a  raw  egg. 
The  albumen  forms  an  insoluble  mass  with  the  poison,  which 
can  then  be  removed  from  the  stomach.  The  common  name 
of  mercuric  chloride  is  corrosive  sublimate.  It  has  strong 
antiseptic  properties  and  is  extensively  used  in  surgery  to 
protect  wounds  from  the  harmful  action  of  germs;  taxider- 
mists sometimes  use  it  to  preserve  skins,  and  it  has  many 
serviceable  applications  as  a  medicine  and  disinfectant.  It 
is  usually  used  as  a  dilute  solution  (1  part  to  1000  parts  of 
water).  Native  mercuric  sulphide  or  cinnabar  (HgS) 
is  a  red  crystalline  solid.  When  hydrogen  sulphide  is  passed 
into  a  solution  of  a  mercuric  salt,  mercuric  sulphide  is 
formed  as  a  black  precipitate;  this  variety,  when  heated, 
changes  into  red  crystals.  Vermilion  is  artificial  mercuric 
.sulphide.  It  is  manufactured  either  (1)  by  grinding  together 
mercury  and  sulphur  and  treating  this  mass  with  potassium 
hydroxide,  or  (2)  by  heating  mercury  and  sulphur  in  iron 
pans  and  subliming  the  black  mass.  In  both  processes  the 
product  must  be  carefully  ground,  washed,  and  dried. 


490  INORGANIC  CHEMISTRY 

Chinese  vermilion  is  the  best  quality.  Vermilion  has  a 
brilliant  red  color,  and,  although  expensive,  is  widely  used 
to  make  red  paint.  Mercurous  nitrate  (HgNO3  or  Hg2(NO3)2) 
and  mercuric  nitrate  (Hg(NO3)2)  are  prepared  by  treating 
mercury  respectively  with  cold  dilute  nitric  acid  and  warm 
concentrated  nitric  acid.  They  are  white,  crystalline  solids. 
Soluble  mercurous  compounds  yield  mercurous  ions  (Hg+) 
and  mercuric  compounds  mercuric  ions  (Hg++).  Both 
kinds  are  colorless.  Mercury  also  forms  complex  ions,  e.g. 
Hg(CN)4— . 

Tests  for  Mercury.  —  Clean  copper  becomes  coated  with  a 
bright  film  of  mercury  when  put  into  the  solution  of  any  mer- 
cury compound.  Other  metals  act  similarly.  (See  Displace- 
ment of  Metals,  Chapter  XXVI.)  Hydrochloric  acid  precipi- 
tates white  mercurous  chloride  from  solutions  of  mercurous 
salts.  Stannous  chloride  in  excess  reduces  mercuric  salts 
at  first  to  white  mercurous  chloride  and  finally  to  a  gray 
precipitate  of  finely  divided  mercury,  thus  :  — 

2HgCl2     +     SnCl2     =     2HgCl     +     SnCl4 

Mercuric  Stannous  Mercurous  Stannic 

Chloride  Chloride  Chloride  Chloride 

2HgCl      +     SnCl2     =       2Hg       +     SnCl4 

Mercurous  Stannous  Mercury  Stannic 

Chloride  Chloride  Chloride 

This  is  a  typical  illustration  of  the  broad  use  of  the  terms 
"oxidation"  and  "reduction,"  viz.  the  addition  and  removal 
of  a  negative  element.  In  these  two  chemical  changes  the 
stannous  chloride  is  oxidized  to  stannic  chloride,  i.e.  the 
negative  element  chlorine  is  added  chemically  to  stannous 
chloride.  On  the  other  hand,  mercuric  and  mercurous 
chloride  both  lose  chlorine,  becoming  respectively  mercu- 
rous chloride  and  mercury,  i.e.  the  negative  element  chlorine 
is  removed  chemically  from  the  two  mercury  compounds. 


ZINC  FAMILY  491 

Mercurous  salts  yield  a  black  precipitate  (e.g.  ammono-basic 
mercuric  chloride  and  mercury  (HgNH2Cl  +  Hg))  with 
ammonium  hydroxide,  while  mercuric  salts  yield  a  white  one 
(e.g.  ammono-basic  mercuric  chloride  (HgNH2Cl)). 

The  Zinc  Family.  —  The  four  elements  just  considered 
(together  with  beryllium)  constitute  a  natural  family  in  the 
periodic  classification.  They  bear  certain  resemblances  to 
the  calcium  family,  which  is  in  the  same  periodic  group,  and. 
the  different  members  resemble  each  other;  but  the  family 
is  not  a  unit,  nor  does  it  exhibit  the  progressive  change  in 
properties  which  characterizes  certain  families.  Zinc  and 
cadmium  are  much  alike,  while  mercury  differs  somewhat 
from  these  metals  and  resembles  copper.  As  already  stated 
these  metals  have  the  valence  two,  except  mercury,  whose 
valence  is  one  in  mercurous  compounds  and  two  in  mercuric 
compounds. 

PROBLEMS  AND  EXERCISES 

1.  How  much  magnesium  will  be  formed  by  heating  100  gm.  of 
potassium    with    magnesium     chloride  ?       (Assume    2  K  +  MgCl2  = 
Mg  +  2KC1.) 

2.  What  is  the  per  cent  of  magnesium  in  (a)  magnesite  (MgCOs), 
(6)  dolomite  (MgCa(CO3)2),  (c)  Epsom  salts  (MgSO4.7  H2O)? 

3.  What  is  the  per  cent  of  zinc  in   (a)  zinc  sulphate  (ZnSO-i), 
(6)  zinc  sulphide  (ZnS),  (c)  zinc  chloride  (ZnCl2),  (d)  zinc  oxide  (ZnO)? 

4.  How  much  zinc  sulphate  can  be  prepared  from  65  gm.  of  zinc? 
from  130  gm.  ?  from  720  gm.  ? 

5.  How  much  mercury  is  formed  by  decomposing  400  gm.  of  cinna- 
bar?    (Assume  HgS  +  O2  =  Hg  +  SO2.) 

6.  What  is  the  per  cent  of  mercury  in  (a)  mercuric  oxide  (HgO), 
(6)  calomel  (HgCl),  (c)  corrosive  sublimate  (HgCl2)? 

7.  What  is  the  approximate  specific  heat  of  cadmium,  accepting 
112  as  its  atomic  weight? 

8  If  64.2501  gm.  of  cadmium  sulphate  yield  44.4491  gm.  of  cad- 
mium sulphide,  what  is  the  atomic  weight  of  cadmium  (assuming 
O  =  16andS  =  32.07)? 


492  INORGANIC  CHEMISTRY 

9.    If  380.5744  gm.  of  mercuric  oxide  yield  352.4079  gm.  of  mer- 
cury, what  is  the  atomic  weight  of  this  element? 

10.  Write  the  formulas  of  the  following  compounds  by  applying 
the  principle  of  valence  and  utilizing  analogous  formulas  in  this 
chapter :     Magnesium  iodide,  ammonium   magnesium  phosphate, 
magnesium  fluoride,   magnesium  silicate,   potassium  zincate,  zinc 
acetate,  zinc  iodide,  cadmium  iodide,  cadmium  sulphate,  mercurous 
bromide,  mercuric  iodide,  mercuric  sulphate. 

11.  What  volume  of  mercury  is  needed  to  fill  one  of  the  commer- 
cial flasks  holding  75  Ib.  ? 

12.  The  freezing  point  of  mercury  is  —  39.5°  C.  and  the  boiling 
point  is  357°  C.     What  are  the  corresponding  Fahrenheit  tempera- 
tures ? 

13.  Give  the  name  and  formula  of  each  magnesium  double  salt 
found  at  Stassfurt  and  indicate  the  ions  found  in  a  dilute  aqueous 
solution  of  each.     (SUGGESTION.  —  See  under  Potassium.) 

14.  What  weight  of   (a)  crystallized  magnesium  sulphate  and 
(6)  crystallized  zinc  sulphate  can  be  made  from  275  gm.  of  the  re- 
spective metals  ? 

15.  What  weight  of  crystallized  zinc  chloride  can  be  made  from 
hydrochloric  acid  and  45  gm.  of  zinc  oxide? 

16.  Calculate  the  simplest  formulas  from  (a)  Hg  =  84.92,  Cl  = 
15.07;   (6)  Hg  =  73.8,  Cl  =  26.2. 

17.  A  solution  of  hydrochloric  acid  contains  39.1  per  cent  of  HC1. 
What  weight  of  zinc  (95  per  cent  pure)  is  needed  to  liberate  the 
hydrogen  from  172  gm.  of  the  solution? 

18.  Calculate  the  solubility  product  of  magnesium  hydroxide 
(Mg(OH)2)  if  the  ionization  is  80  per  cent  and  the  molar  solubility 
is  .00015.     And  of  mercuric  sulphide  (HgS)  if  the  ionization  is  100 
per  cent  and  the  molar  solubility  is  .02el7. 


CHAPTER  XXIX 
Aluminium 

Occurrence. — Aluminium  (or  aluminum),  Al,  does  not 
occur  free  in  nature,  but  its  compounds  are  numerous  and 
widely  distributed.  About  8  per  cent  of  the  earth's  crust 
is  aluminium.  It  is  the  most  abundant  metal  and  the  third 
element  in  abundance  in  the  earth's  crust.  Many  common 
rocks  are  silicates  of  aluminium  and  other  metals,  e.g.  feldspar 
and  mica,  which  make  up  a  large  part  of  granite  and  gneiss. 
Clay  and  slate  are  mainly  silicate  of  aluminium,  which  was 
formed  by  the  decomposition  of  complex  aluminium  miner- 
als. Corundum  and  emery  are  more  or  less  impure  alumin- 
ium oxide  (A12O3).  Bauxite  is  an  impure  hydroxide  of  alu- 
minium (A12O5H4).  Cryolite  is  a  fluoride  of  aluminium  and 
sodium  (Na3AlF6). 

Metallurgy.  — Aluminium  is  obtained  from  its  purified 
oxide  (A12O8)  by  electrolysis.  An  open  iron  vessel  lined  with 
carbon  is  made  the  cathode  (Fig.  77).  The  anode  consists 
of  several  carbon  bars  hung  from  a  common  copper  rod, 
which  can  be  lowered  as  the  carbon  is  consumed.  The  bot- 
tom of  the  box  is  first  covered  with  cryolite,  the  anode  is 
lowered,  and  the  box  is  then  filled  with  cryolite.  The  cur- 
rent is  turned  on,  and  in  its  resisted  passage  through  the 
cryolite  enough  heat  is  generated  to  melt  the  cryolite.  Pure, 
dry  aluminium  oxide  is  now  added,  which  is  decomposed 
into  aluminium  and  oxygen.  The  oxygen  goes  to  the  anode 
and  unites  with  the  carbon,  forming  carbon  monoxide,  which 
burns  or  escapes.  The  molten  aluminium,  which  goes  to  the 

493 


494 


INORGANIC   CHEMISTRY 


cathode,  gradually  collects  on  the  bottom  of  the  cell.     The 
process  is  continuous,  fresh  aluminium  oxide  being  added 

and  the  molten  alumin- 
ium being  drawn  off  at 
intervals.  The  cryolite  is 
unchanged  chemically. 


D 


A 


Properties.  — Alumin- 
ium is  a  lustrous,  bluish 
white  metal.  It  is  very 
light  compared  with 
other  common  metals, 
since  its  specific  gravity 


FIG.  77.  — Apparatus  for  the  manufacture  of 
aluminium  by  the  electrolysis  of  aluminium 
oxide.  C,  C,  C  is  the  iron  box  lined  with 
carbon  which  serves  as  a  cathode.  A,  A, 
etc.,  are  carbon  bars  attached  to  the  cop- 
per rod  R.  Connection  is  made  with  the  is  only  about  2.6;  this 
cathode  at  D.  ,  . ,  .  ,  , , 

value  is  one  third  that 

of  iron.  It  is  ductile  and  malleable,  and  is  often  sold 
in  the  form  of  wire  and  sheets;  it  must  be  annealed  fre- 
quently during  the  hammering  or  drawing.  It  is  a  good 
conductor  of  heat  and  electricity.  Its  tensile  strength  is 
about  as  great  as  that  of  cast  iron.  It  melts  at  about  658° 
C.  and  can  be  cast  and  welded,  but  not  readily  soldered  so  as 
to  produce  a  permanent  joint.  Pure  aluminium  is  only  very 
slightly  oxidized  by  air.  Hydrochloric  acid  interacts  readily 
with  it,  forming  aluminium  chloride  and  hydrogen,  thus :  — 


2A1 

Aluminium 


6HC1 

Hydrochloric 
Acid 


2A1C13 

Aluminium 

Chloride 


3H2 

Hydrogen 


Sulphuric  acid  interacts  feebly  with  aluminium,  but  nitric 
acid  has  no  apparent  effect.  With  a  boiling  solution  of 
sodium  or  potassium  hydroxide  aluminium  forms  hydrogen 
and  an  aluminate,  thus  :  — 


6NaOH        +        2A1       = 

Sodium  Hydroxide  Aluminium 


2  Na3AlO3 

Sodium 
Aluminate 


3H2 

Hydrogen 


ALUMINIUM  495 

Uses.  —  The  varied  properties  of  aluminium  adapt  it  to 
numerous  uses.  It  is  made  into  the  metallic  parts  of  military 
outfits,  caps  for  jars,  surgical  instruments,  cooking  utensils, 
tubes,  the  fittings  of  boats,  automobiles,  and  air  ships,  tele- 
phone receivers,  scientific  apparatus,  parts  of  opera  glasses 
and  telescopes,  the  framework  of  cameras,  stock  patterns  for 
foundry  work,  and  hardware  samples.  Its  attractive  appear- 
ance has  led  to  its  extensive  use  as  an  ornamental  metal,  both 
in  interior  decorative  work  and  in  numerous  small  objects, 
such  as  trays,  picture  frames,  hairpins,  and  combs.  Alumin- 
ium leaf  is  used  for  decorating  book  covers  and  signs;  a 
mixture  of  aluminium  powder  and  an  adhesive  oil  is  likewise 
used  as  a  paint  for  steam  pipes,  lamp-posts,  radiators,  smoke- 
stacks, and  other  metal  objects.  Aluminium  wire  is  used 
as  a  conductor  of  electricity.  Large  quantities  of  alumin- 
ium are  used  in  the  steel  industry;  added  to  molten  steel, 
the  aluminium  combines  with  gases  and  produces  castings 
free  from  blow  holes.  Considerable  is  used  in  preparing  cer- 
tain metals  and  in  welding.  Thus,  if  a  mixture  of  chromium 
oxide  and  powdered  aluminium  is  ignited  at  one  point  by 
a  special  device,  the  reduction  thus  initiated  proceeds  rapidly 
throughout  the  mixture  and  the  intense  heat  thereby  gener- 
ated fuses  the  chromium,  which  can  be  removed  from  the 
crucible  subsequently  as  a  coherent  mass;  the  aluminium 
oxide  likewise  melts  and  separates  from  the  metal.  The 
equation  for  the  chemical  change  is  — 

Cr2O3     +     2A1     =       2Cr      +      A12O3 

Chromium        Aluminium         Chromium  Aluminium 

Oxide  Oxide 

Other  metals  hitherto  rare  or  expensive  can  be  similarly  pre- 
pared. If  a  mixture  of  ferric  oxide  (Fe2O3)  and  powdered 
aluminium  is  ignited,  molten  steel  at  a  temperature  of  about 
3000°  C.  is  produced.  By  means  of  a  special  apparatus  the 
molten  steel  can  be  conducted  to  a  joint  or  crack ;  this  pro- 


496  INORGANIC  CHEMISTRY 

cess  is  used  in  welding  iron  rails  and  repairing  fractures. 
These  mixtures  of  aluminium  and  oxides  are  called  "  thermit," 
and  the  method  is  known  as  the  Goldschmidt  or  alumino- 
thermic  method. 

Alloys.  —  The  alloy  of  aluminium  and  copper  —  alumin- 
ium bronze  —  has  been  described.  (See  Alloys  of  Copper.) 
Magnalium  contains  from  75  to  90  per  cent  of  aluminium,  the 
rest  being  magnesium;  it  is  used  in  scientific  instruments, 
e.g.  as  balance  beams. 

Aluminium  Oxide,  or  alumina,  A12O3,  is  the  only  oxide  of 
aluminium.  Pure  crystalline  varieties  of  native  alumina  are 
known  as  corundum,  while  the  impure  forms  are  called 
emery ;  they  are  very  hard  substances,  corundum  being  next 
below  the  diamond  in  the  scale  of  hardness.  Emery  was 
formerly  used  as  an  abrasive,  but  it  has  been  largely  displaced 
by  carborundum  (see  page  385).  The  transparent  colored 
kinds  of  corundum  are  prized  as  gems  (see  below). 

Alumina  is  prepared  as  a  white  powder  by  heating  the 
metal  in  the  air  or  by  heating  the  hydroxide.  The  product 
obtained  by  heating  the  hydroxide  to  a  low  temperature 
interacts  with  acids  but  the  intensely  heated  oxide  resembles 
the  crystallized  varieties  in  being  almost  insoluble  in  acids. 
The  native  oxide  is  converted  into  an  aluminate  and  a  sul- 
phate respectively  by  fusing  with  a  caustic  alkali  (e.g.  potas- 
sium hydroxide)  and  with  acid  potassium  sulphate  (HKSO4). 
Alumina  melts  in  the  oxy hydrogen  flame,  in  the  electric  fur- 
nace, and  during  the  reduction  of  oxides  by  aluminium  (see 
Goldschmidt  method  above). 

Compact  crystalline  alumina  is  manufactured  from  bauxite 
by  an  electrothermic  process.  It  resembles  native  corundum 
and  is  used  as  an  abrasive.  This  artificial  alumina  is  known 
in  trade  as  "alundum." 


ALUMINIUM  497 

Gems  containing  Aluminium.  —  Corundum  (A1203)  has  long  been 
found  in  crystals  in  Ceylon,  Siam,  Burma,  and  other  places  in  the 
Orient.  The  color  is  due  to  traces  of  impurities,  usually  oxides  of 
metals.  The  sapphire  is  blue  and  the  ruby  is  red.  The  Oriental 
topaz  is  yellow,  the  Oriental  amethyst  is  purple,  and  the  Oriental 
emerald  is  green.  Sapphires,  rubies,  and  similar  gems  are  now  made 
by  melting  aluminium  oxide  (with  or  without  coloring  matter)  in  an 
oxy hydrogen  flame.  Spinels  are  complex  compounds  of  alumin- 
ium. The  typical  or  ruby  spinel  is  magnesium  alumina te  (MgAl2O4). 
It  resembles  the  true  ruby  in  color.  Other  spinels  differ  from 
the  ruby  spinel  both  in  color  and  in  composition.  Turquoise  is 
a  complex  aluminium  phosphate  containing  traces  of  copper.  It 
has  a  beautiful  robin's-egg-blue  color,  is  compact,  and  may  be  worked 
into  various  shapes.  Formerly  turquoise  came  almost  exclusively 
from  Persia,  but  now  New  Mexico  meets  most  demands.  Topaz  is 
a  complex  aluminium  silicate  containing  fluorine,  usually  has  a  pale 
yellow  color,  and  is  found  in  many  localities.  Emerald  is  the  most 
precious  gem  next  to  diamond  and  ruby.  It  is  an  aluminium  silicate 
containing  the  rare  element  beryllium.  The  finest  specimens  have  a 
deep  emerald  green  color  and  come  from  Colombia,  South  America. 
Garnet  is  a  complex  silicate  of  aluminium  and  another  metal,  espe- 
cially calcium,  magnesium,  iron,  or  manganese.  The  kind  used  as  a 
gem  has  a  deep  red  color  and  is  rather  abundant. 

Aluminium  Hydroxide,  A1(OH)3,  is  a  white,  jelly  like  solid 
formed  by  the  interaction  of  an  hydroxide  and  the  solution 
of  an  aluminium  salt,  e.g. :  — 

A1C13    +     3NH4OH     =     A1(OH)3     +     3NH4C1 

Aluminium  Ammonium  Aluminium  Ammonium 

Chloride  Hydroxide  Hydroxide  Chloride 

It  has  weak  acid  and  basic  properties,  the  latter,  however, 
being  the  stronger.  Its  acid  property  is  shown  by  the  forma- 
tion of  soluble  saltlike  compounds  called  aluminates  upon 
the  addition  of  an  excess  of  sodium  or  potassium  hydroxide, 
thus :  — 

A1(OH)3      +      SNaOH      =      Na3AlO3      +     3H2O 

Aluminium  Sodium  Sodium  Water 

Hydroxide  Hydroxide  Aluminate 


-jus  INORGANIC  UII<;.MISTI;Y 

The  feeble  basic  properly  <>!'  :i  In  ininiu  in  hydwxido  is  illus- 
trated |,y  I  IK-  fact  that  it  docs  not  form  s:ill.s  with  such  \\<-;ik 
acids  as  carbonic,  1  1\  «  I  n  isu  1  1  iliu  rir  (ILS),  :iiul  sulphu  n  Ml:; 
(I  LS(  )..,).  Tims,  \\lien  sodium  carbonate  or  ;i  iiinioniiini 

;;i|||>liidr   is    :nM(><|    hi    I  lie   solution    of   ;in    aluminium    salt,  alil- 

hydroxide   (not    (lie  c.-irhonah'  or  sulphide)   is  |uv- 
,  e.g.:  - 


f  *  (NHOiS  -f  6  HjO  -  2  A1(OH)8  +  6  NH4CI  4  :\  I  I..S 

\lniiiin  Ammo  \\;il<i  Alumiii-  Ammo  ll\<li.. 

IMIM  Ililim  MINI  I  .......  !•(•!! 

Clil..ii.|n  Hiilpliltlu  My<lro\i.l«-  ri.l..ri.l.-       Sul|.lii.lo 

Like  nil  hnscs,  lio\\c\  cr,  ;i  In  mi  inn  in  h\«lro\'nlr  inleracls  with 
strong  arids  and  Ilicrch     forms  ;  ;i  ll    ,  llius: 


Al(OH)a       +       :UI(M  A1(113       4-       3H80 

Allllllillilim  I  ly.lrochloiic  \  IUHMHIIIIII  \\nliT 

Helios!,!.'  \.M|  Chi  .....  I- 

Commrrrial  a  In  mi  inn  in  hydroxide  is  j  m-pa  red  I  iy  n»:isl  in;;- 
l>aii\ile  or  eryolih-  \\ilh  sodium  ca  rhoiial  e,  e\  I  rac!  i  ni;  i.lu» 
resulting  iduminale  \\ilh  \\ah-r,  and  precipitating  the  hy- 
droxide l>y  IIMSS'UI!';  carhoii  dioxide  into  I  he  solution;  the 
last  operat  ion  may  l>e  rc|iresenl  ed  l>\'  I  he  e<|iia  t  ion  : 

2Na8Al()a  4-    3  CO,    4-  -'nia()  =  L>Al(()ll),  -f-  :;\a.co; 

S.Mliinu  C:II|M,II  \\  :I!IT  A  linn  mi  urn  S.i.lmm 

Aiu  .....  L.I.  i>i,.M,ir  HydroaticU 


The  aluminium  hydroxide  is  dried  and  sold  in  the  form  of  u 
\\hile  powder.  There  are  several  native  alnmin'mm  hydrox 
idos.  Bauxite  (Al,(  )fll  I4  or  Al,():l  .  L»  I  !.,())  contains  ferric  oxide 
i  I  .  <»  MS  an  impurity.  It  rcscmhlcs  clay  in  texture  and 
color.  The  \asl  deposits  found  at  I'.anx.  in  southern  Trance, 
I'urnis.h  much  of  I  lie  raw  material  for  the  manufacture  of 
aluminium,  though  some  is  obtained  from  our  Southern 
states.  llydran-vllito  (Al(Oll).,)  and  diasporo  (AlOJl)  aiv 
found  in  relatively  small  <|iianli(ics. 


ALUMINIUM  I!)!) 

Aluminates  have  been  described  in  the  foregoing  para- 
graphs. They  are  compounds  in  \\hich  aluminium  acts  :is 
:in  acid  element,  corresponding  in  this  respeet  to  the  /incales. 
Aluminates  are  soluble  to  sonic  extent  in  \v:iter  ;ind  such 
solutions  have  :in  :dk:dine  reaction  owiii!1;  to  hydrolysis;  tll(5 
equation  for  the  hydrolysis  is 

Na3AK)a  +  3  H,0  •  A1(OH)8  4-  3  NaOH 

tin*  alkaline  reaction  being  caused  by  the  hydroxyl  ions  lib- 
erated hy  the  ioni/;ition  of  the  sodium  hydroxide.  No 

:ilmmn;ile  is  formed  by  adding  an  excess  of  ammonium  hy- 
droxide* to  jduininiiim  hydroxide;  this  property  is  some- 
times used  to  distinguish  :iluminium  from  /inc.  since  //me 
hydroxide  (a  siiuihr  (-(.nipoiind)  forms  :i  soluble  complex 
compound  hy  the  interact  ion  with  ammonium  hydroxide. 

Aluminium  Sulphate,  AI,(S< ).,);, .  ISIlJ),  is  a,  wliit.e,  crystal 
line  solid.  The  commercial  salt  has  a  va.riahle  composition; 
if  purr,  it  dissolves  readily  mid  complotdy  in  wator.  It  is 
extensively  used  as  a  mordant  in  dyeing  (se<>  below),  US  an 
ingredient  of  the  si/e  put,  upon  paper  to  prevent  the  ink  from 
spreading,  and  in  purifying  water. 

A  solution  of  aluminium  sulphate  has  an  acid  reaction  on 
aecount  of  hydrolysis;  the  equation  for  the  hydrolysis  is — 

Ala(SO4)3  +  0  Ha()  =  2  Al(OH),  +  3  I^S()4 
the  arid  reaction  heing  due  to  the  hydrogen  ions  liberated 
hy  the  ioni/ation  of  tlie  sulphui'ic  acid. 

Aluminium  sulphate  is  prepared  from  pure  rl:iy,  l)au\ile,  or 
rrv«.lile.  II  clny  nr  l>:iu\ile  i.;  lienled  with  sulphuric  ;icid  :iu«l  then 
:ill«.\\r(l  Id  cool,  the  product,  is  impure  aluminium  sulphate,  known  as 
":ilnm  cake,"  or  if  much  inm  is  |)resent,  HH  "  illumino  ferric  c;ike";  it 
is  used  |.o  purify  scwu^e  and  lor  oilier  purposes  where  iron  and  MM- 
other  impurities  are  harmless.  Purer  siluiitimum  sulphaie  is  pre|»:n<-d 
hy  hea,tin«  bauxite  \\ilh  s.Hlium  ciirbonnte,  exlraclin^  the  sodium 
aliimiii:ife  with  water,  and  preripitutiiiK  the  .'iluminium  us  the  liyilrox- 


500  INORGANIC   CHEMISTRY 

ide  with  carbon  dioxide  gas;  the  relatively  pure  hydroxide  is  then 
changed  into  sulphate  by  treatment  with  sulphuric  acid.  The 
product,  known  as  "concentrated  alum,"  has  the  composition  ex< 
pressed  by  the  formula  Ala  (804)3  . 20  H^O,  though  separate  crystals 
may  contain  only  eighteen  molecules  of  water  of  crystallization.  By 
boiling  cryolite  with  milk  of  lime,  the  sodium  aluminate  thereby 
formed  may  be  changed  into  "  concentrated  alum,"  as  described  above. 
About  50,000  tons  of  "concentrated  alum"  are  annually  produced 
in  the  United  States. 

Alum.  —  When  solutions  of  aluminium  sulphate  and  po- 
tassium sulphate  are  mixed  and  concentrated  by  evapora- 
tion, transparent,  colorless,  glassy  crystals  are  deposited. 
This  solid  is  potassium  alum  or  simply  alum.  It  has  the 
composition  represented  by  the  formula,  K2A12(SO4)4 .  24  H2O,, 
or  K2SO4,  A12(SO4)3 .  24  H20,  and  is  sometimes  called  a  double 
salt.  It  is  the  type  of  a  class  of  similar  salts  called  alums, 
which  can  be  formed  by  mixing  the  solution  of  a  sulphate  of 
a  trivalent  metal  (e.g.  Al,  Cr,  Fe)  with  the  solution  of  a  sul- 
phate of  a  univalent  metal  (e.g.  K,  Na,  NH4).  Alums  are 
rather  soluble  in  water,  and  their  solutions  have  an  acid 
reaction  owing  to  hydrolysis.  (See  Aluminium  Sulphate.) 
They  crystallize  as  octahedrons  and  contain  twenty-four 
molecules  of  water  of  crystallization.  When  heated,  alums 
lose  their  water  of  crystallization  and  some  sulphur  trioxide 
and  fall  to  a  white  powder  or  porous  mass  known  as  burnt 
alum.  Potassium  alum  is  the  most  common,  but  ammonium 
and  sodium  alums  are  manufactured  and  used.  Sodium 
alum  is  an  ingredient  of  some  baking  powders.  Burnt  alum 
finds  application  as  a  medicine.  Alum  has  been  largely  dis- 
placed by  "  concentrated  alum,"  since  aluminium  sulphate 
has  the  same  general  properties;  but  the  real  alum  is  still 
used  to  some  extent  in  dyeing  and  printing  cloth,  in  tanning 
and  paper  making,  in  purifying  water  and  sewage,  as  a  medi- 
cine, for  hardening  plaster,  in  making  wood  and  cloth  fire- 
proof, and  in  preparing  other  aluminium  compounds. 


ALUMINIUM  501 

Alum  was  known  to  the  ancients,  who  used  it  in  dyeing  and  tan- 
ning, and  as  a  medicine.  It  was  first  manufactured  in  Europe,  about 
the  thirteenth  century,  from  native  alunite,  which  is  an  impure  sul- 
phate of  aluminium,  potassium,  and  iron.  Alunite  and  alum  slates 
or  shales  are  now  used  to  some  extent,  but  most  of  the  alum  is  made 
from  bauxite. 

Not  all  alums  contain  aluminium.  As  stated  in  the  preceding 
paragraph,  this  metal  may  be  replaced  by  iron,  chromium,  manganese, 
or  any  metal  having  the  valence  of  three.  Hence  the  general  formula 
of  an  alum  is  M2  (864)3  .  X2SO4  .  24  H2O,  in  which  M  may  be  aluminium, 
iron,  chromium,  etc.,  and  X  a  metal  (or  group)  like  potassium,  sodium, 
ammonium. 

Alums  are  double  salts,  i.e.  crystalline  compounds  of  two  or  more 
normal  salts.  Their  dilute  aqueous  solutions  contain  the  ions  of  the 
separate  salts  and  have  properties  which  indicate  the  double  nature, 
so  to  speak,  of  the  salt  in  question.  (Compare  the  Stassfurt  salts.) 

Alums  and  other  aluminium  salts  are  used  as  mordants  in 
dyeing  and  calico  printing.  Some  dyes  must  be  fixed  in  the 
fabric  by  a  metallic  substance,  otherwise  the  color  could  be 
easily  removed.  The  cloth  to  be  dyed  or  printed  is  impreg- 
nated or  printed  with  a  mordant,  and  then  heated  or  treated 
with  some  substance  to  change  the  mordant  into  an  insoluble 
compound.  The  mordanted  cloth  is  next  passed  through  a 
vat  containing  the  solution  of  the  dye,  which  unites  chemi- 
cally or  mechanically  (perhaps  both)  with  the  metallic 
compound,  forming  a  colored  compound.  The  latter  is 
called  a  "lake";  it  is  relatively  insoluble,  and  cannot  be 
easily  washed  from  the  cloth,  i.e.  it  is  a  fast  color.  Alumin- 
ium acetate  or  "red  liquor,"  aluminium  sulphate,  and 
sodium  aluminate,  besides  alum,  are  used  as  mordants  for 
cotton,  linen,  and  wool.  The  use  of  aluminium  salts  as 
mordants  depends  upon  the  fact  that  they  hydrolyze  readily; 
it  is  the  resulting  aluminium  hydroxide,  therefore,  which  is 
the  effective  metallic  compound  in  dyeing. 

Cryolite  is  a  white,  glassy,  crystalline  solid.  It  often 
resembles  clouded  ice,  and  its  name  means  "ice  stone." 


502  INORGANIC  CHEMISTRY 

Its  composition  corresponds  to  the  formula  Na3AlF6  (or 
A1F3.  3  NaF).  Small  fragments  melt  easily,  even  in  a  candle 
flame,  and  color  the  Bunsen  flame  yellow.  The  only  local- 
ity where  it  is  found  in  commercial  quantities  is  southern 
Greenland,  which  yields  annually  about  10,000  tons.  It  is 
used  not  only  in  manufacturing  aluminium,  but  as  a  source 
of  alum  and  aluminium  hydroxide,  pure  sodium  carbonate 
and  hydroxide,  hydrofluoric  acid,  and  fluorides. 

Aluminium  Chloride  when  pure  is  a  white  powder,  but  it  is 
often  a  yellowish,  crystalline  mass  (A1C13 .  6  H2O).  It  is  pre- 
pared by  heating  powdered  aluminium  in  chlorine,  or  by 
passing  chlorine  over  a  heated  mixture  of  aluminium  oxide 
and  carbon.  Exposed  to  the  air,  it  absorbs  moisture  and 
gives  off  fumes  of  hydrochloric  acid.  It  dissolves  in  water 
with  evolution  of  heat,  and  if  the  solution  is  heated,  hydro- 
chloric acid  is  expelled,  owing  to  the  hydrolysis  of  the  chloride, 
thus : — 

A1C13     +     3H2O     =      3HC1     +     A1(OH)3 

Aluminium  Water  Hydrochloric         Aluminium 

Chloride  Acid  Hydroxide 

This  salt  is  used  in  organic  chemistry. 

Tests  for  Aluminium.  —  When  a  compound  of  aluminium 
is  heated  on  charcoal  with  a  blowpipe,  cooled,  moistened 
with  cobaltous  nitrate  solution,  and  then  reheated,  the  mass 
finally  turns  a  beautiful  blue  color.  Sodium  hydroxide  pre- 
cipitates white  gelatinous  aluminium  hydroxide,  which  dis- 
solves in  an  excess  of  the  alkali  (NaOH),  owing  to  the  forma- 
tion of  soluble  sodium  aluminate.  Aluminium  hydroxide 
is  insoluble  in  an  excess  of  ammonium  hydroxide  (distinc- 
tion from  zinc  hydroxide). 

Miscellaneous.  —  It  is  evident  from  the  preceding  para- 
graphs that  aluminium,  like  zinc,  acts  both  as  a  metal  and  a 


ALUMINIUM  503 

non-metal.     Many  soluble  aluminium  compounds  yield  color- 
less aluminium  ions  (A1+  +  +). 

The  atomic  weight  of  this  element  is  27.1  and  the  valence 
is  three. 

Clay  is  a  more  or  less  impure  aluminium  silicate,  formed  by  the 
slow  decomposition  of  rocks  containing  aluminium,  especially  feld- 
spar. Pure  feldspar  is  a  silicate  of  aluminium  and  sodium  or  potas- 
sium. The  products  of  its  decomposition  are  chiefly  an  insoluble 
aluminium  silicate  and  a  soluble  alkaline  silicate.  The  latter  is 
washed  away.  The  aluminium  silicate  which  remains  is  pure  clay 
or  kaolin.  The  latter  is  really  a  hydrous  silicate,  having  the  com- 
position corresponding  to  the  formula  A^SiaOy .  2  H2O.  The  com- 
position of  clay  varies,  because  it  is  seldom  formed  from  pure  feldspar. 
Pure  kaolin  is  a  white  powdery  substance,  but  most  kaolin  contains 
particles  of  mica  and  quartz.  Ordinary  clay  contains  many  impuri- 
ties, e.g.  carbonates  of  calcium  and  magnesium,  quartz,  and  iron 
compounds.  All  kinds  of  clay  become  plastic  when  wet  and  can  be 
molded  into  various  objects  which  shrink  on  drying  but  retain  their 
general  form;  if  heated,  the  dried  clay  does  not  melt  (except  at  a 
very  high  temperature)  but  becomes  a  permanently  hard  mass. 
These  properties  (plasticity  when  wet  and  hardness  when  heated) 
make  clay  a  most  serviceable  substance.  For  ages  it  has  been  made 
into  useful  and  ornamental  articles  which  may  be  roughly  put  into 
three  comprehensive  classes  —  porcelain,  pottery,  and  materials  of 
construction.  The  classes  as  well  as  the  varieties  in  each  class  are 
the  result  of  differences  in  quality  and  proportions  of  raw  material 
and  in  method  of  heating.  The  subvarieties  merge  into  each  other. 

Porcelain  (china  or  chinaware)  is  the  finest  clay  product.  It  is 
made  by  fusing  a  mixture  of  very  pure  kaolin,  fine  sand,  and  a  more 
fusible  substance,  usually  feldspar,  though  sometimes  chalk  and 
gypsum  are  also  used.  The  fused  mass  when  cool  is  hard,  dense, 
white,  and  translucent  (if  thin);  it  is  often. called  "biscuit"  or  "bis- 
cuit ware."  Although  not  very  porous,  its  surface  is  glazed,  partly 
for  protection,  partly  for  ornament.  This  is  done  by  coating  the 
ware  with  a  thin  mixture  similar  to  that  used  for  making  the  porcelain 
but  more  easily  fused,  and  then  heating  again  so  that  the  glaze  will 
penetrate  the  clay.  Pottery  is  a  very  large  class  and  includes  colored, 
white,  glazed,  and  unglazed  ware.  The  raw  material  is  not  as  pure 


504  INORGANIC  CHEMISTRY 

as  that  used  for  porcelain,  nor  is  the  mixture  heated  to  such  a  high 
temperature;  the  product,  therefore,  is  rather  coarse,  opaque,  heavy, 
and  porous.  The  finer  varieties,  such  as  crockery,  art  pottery,  and 
some  kinds  of  stone  ware,  are  glazed  like  porcelain.  The  coarser 
varieties,  such  as  jugs  and  domestic  utensils,  are  glazed  by  throwing 
salt  into  the  oven  just  before  the  firing  (i.e.  the  baking  or  heating)  is 
over;  the  sodium  in  the  sodium  chloride  forms  a  fusible  aluminium 
silicate,  which  coats  the  surface.  Unglazed  pottery  is  familiar  under 
different  names,  e.g.  flower  pots,  tiles,  terra  cotta,  and  clay  tobacco 
pipes.  Materials  of  construction  made  from  clay  include  some  grades 
of  stoneware,  bricks,  pipes,  drain  tile,  etc.  The  raw  material  is  usu- 
ally impure  and  the  firing  is  done  at  a  low  temperature.  The  product 
varies  with  the  quality  of  the  clay.  Thus,  ordinary  clay  containing 
iron  compounds  gives  coarse  red  brick,  while  clay  containing  consider- 
able silica  gives  bricks  which  withstand  high  temperature  and  are 
called  fire-clay  bricks.  Pipe  and  drain  tile,  both  glazed  and  porous, 
are  used  to  convey  water,  sewage,  and  other  fluids,  and  as  a  conduit 
for  underground  electric  wires,  especially  wire  cables. 

PROBLEMS  AND  EXERCISES 

1.  What  is  the  per  cent  of  aluminium  in  (a)  cryolite  (NasAlF6), 
(6)    turquoise  (A12P2O8  .  H6A12O6 . 2  H20),  (c)    corundum  (A12O3),  (d) 
aluminium  hydroxide? 

2.  What  volume  of  oxygen  at  15°  C.  and  760  mm.  is  needed  to 
change  5  gm.  of  aluminium  to  aluminium  oxide  (Al2Oa)  ? 

3.  If  6.917   gm.  of    aluminium    bromide   (AlBr3)  require  8.4429 
gm.  of  silver  to  precipitate  all  the  bromine,  what  is  the  atomic  weight 
of  aluminium?     (Assume  Ag  =  107.88  and  Br  =  79.92.) 

4.  Write  the  formulas  of  the  following  compounds  by  applying 
the  principle  of  valence  and  calculate  the  per  cent   of  aluminium  in 
each:  Aluminium  sulphide,  aluminium  phosphate,  aluminium  acetate, 
potassium  aluminate. 

5.  Compare  the  corresponding  compounds  of  aluminium  and 
zinc. 

6.  Discuss  (a)  "  aluminium  acts  as  a  metal  and  a  non-metal," 
and  (6)  hydrolysis  of  aluminium  compounds. 


CHAPTER  XXX 
Tin  and  Lead  —  Cerium  and  Thorium 

TIN  and  lead  are  familiar  metals.  They  have  similar  and 
useful  properties,  which  give  these  metals  and  their  com- 
pounds numerous  applications. 

TIN 

Occurrence  of  Tin.  —  Metallic  tin  is  rarely  if  ever  found. 
Tin  dioxide  (cassiterite  or  tin  stone,  SnO2)  is  the  only  avail- 
able ore.  It  is  not  widely  distributed,  but  large  deposits 
are  found  in  England  (at  Cornwall),  Germany  (in  Bohemia 
and  Saxony),  Australia,  Tasmania,  and  the  East  Indian 
Islands,  especially  Banca  and  Billiton.  A  small  quantity  is 
found,  but  not  mined,  in  the  United  States. 

Tin  is  one  of  the  oldest  known  metals.  It  is  mentioned  in  the 
Pentateuch  and  was  probably  obtained  long  before  the  Christian 
era  by  the  Phoenicians  from  the  British  Isles,  which  were  called 
Cassiterides  (from  the  Greek  word  kassiteros,  meaning  tin).  Many 
ancient  bronzea  contain  tin.  The  alchemists  called  it  Jupiter,  and 
used  the  metal  and  its  compounds. 

The  Latin  word  stannum  gives  the  symbol  Sn  and  the  terms 
stannous  and  stannic. 

Metallurgy  of  Tin.  —  If  the  tin  ore  contains  sulphur  or 
arsenic,  these  impurities  must  be  removed  by  roasting. 
The  tin  oxide  is  then  reduced  by  heating  it  with  coal  in  a 
reverberatory  furnace;  the  simplest  equation  for  this  change 
is  — 

SnO2     +     C     =     Sn     +     CO2 

Tin  Carbon  Tin  Carbon 

Dioxide  Dioxide 

505 


506  INORGANIC  CHEMISTRY 

The  molten  tin,  which  collects  at  the  bottom  of  the  furnace, 
is  drawn  off  and  cast  into  bars  or  masses,  which  are  often 
called  block  tin.  Usually  it  is  purified  by  melting  it  slowly 
on  a  hearth,  inclined  so  that  the  more  easily  melted  tin  will 
flow  down  the  hearth  and  leave  the  metallic  impurities  be- 
hind. This  tin  may  be  further  purified  by  stirring  the  molten 
metal  with  a  wooden  pole  or  by  holding  billets  of  wood 
beneath  its  surface.  The  impurities,  which  are  oxidized  by 
the  escaping  gases,  collect  as  a  scum  on  the  surface  and  are 
removed. 

A  small  amount  of  tin  is  obtained  by  treating  rejected 
tin  plate  or  scrap  tin  with  chlorine  or  some  other  active 
dissolving  chemical. 

Properties  of  Tin.  —  Tin  is  a  white,  lustrous  metal.  It 
is  soft  and  malleable,  and  can  be  readily  cut  and  hammered. 
It  is  softer  than  zinc  but  harder  than  lead.  Its  specific 
gravity  is  7.3.  Tin  may  be  obtained  in  the  crystalline 
form,  and  when  a  piece  of  such  tin  is  bent  it  makes  a  crac- 
kling sound,  which  is  caused  by  the  friction  of  these  crystals 
upon  one  another.  It  melts  at  about  232°  C.,  and  when 
heated  to  a  higher  temperature  it  burns,  forming  white 
tin  oxide  (SnO2).  The  physical  properties  of  tin,  like  those 
of  zinc,  vary  with  the  temperature.  One  property  of 
tin  is  rather  striking.  If  kept  at  a  low  temperature  for 
some  time,  ordinary  tin  changes  slowly  into  a  gray  powder 
having  a  specific  gravity  of  about  5.8.  Experiment  shows 
that  this  transformation  begins  at  18°  C.,  and  this  tempera- 
ture is  called  the  transition  point.  That  is,  ordinary  tin  is 
stable  only  above  18°  C. ;  below  18°  C.  it  is  unstable  and  may 
form  gray  tin.  Sometimes  this  "  tin  disease,"  as  it  might  be 
called,  attacks  the  pipes  of  church  organs.  The  appearance 
of  a  sheet  of  tin  affected  by  "  tin  disease  "  is  shown  in  Fig- 
ure 78.  Concentrated  hydrochloric  acid  changes  tin  into 


TIN  507 

stannous  chloride  (SnCl2),  hot  concentrated  sulphuric  acid 
converts  it  into  stannous  sulphate  (SnSO4),  while  commer- 
cial nitric  acid  transforms  it  into  a  white  solid  (metastannic 
acid  (H2Sn03)5).  With  sodium  hydroxide  tin  forms  sodium 


FIG.  78.  —  Sheet  of  tin  affected  by  "  tin  disease  "  (enlarged  one  and  one 
half  times). 

metastannate  (Na2SnO3).  Zinc  precipitates  tin  from  its  so- 
lutions as  a  grayish  black,  spongy  mass,  which  is  sometimes 
filled  with  bright  scales.  This  is  due  to  the  fact  that  zinc 
precedes  tin  (by  several  places)  in  the  electromotive  series 
of  the  metals.  (See  Chapter  XXVI,  Electromotive  Series  of 
Metals.) 

Uses  of  Tin.  —  Tin  is  so  permanent  in  air,  water,  weak 
acids  (like  vinegar  and  fruit  acids),  and  alkalies  that  it  is 
extensively  used  in  making  scientific  apparatus  and  as  a 
protective  coating  for  metals.  Condensing  pipes  in  stills 
are  often  made  of  tin.  Ordinary  tinware  is  sheet  iron 
coated  with  tin.  The  tin  plate  (sheet  tin,  or  simply  "tin") 
is  made  by  dipping  very  clean  sheet  iron  into  molten  tin. 
Tacks,  nails,  and  many  small  iron  objects  are  similarly 
tinned.  Copper  coated  with  tin  is  made  into  vessels  for 
cooking,  and  brass  coated  with  tin  is  made  into  pins.  Large 


508  INORGANIC  CHEMISTRY 

quantities  of  tin  plate  are  used  to  cover  roofs.  Tinned  iron 
does  not  rust  until  the  tin  is  worn  off  and  the  iron  exposed, 
and  then  the  rusting  proceeds  rapidly.  Tin  is  also  ham- 
mered into  thin  sheets  called  tin  foil,  though  much  tin  foil 
contains  lead.  Many  useful  alloys  contain  tin  as  an  essen- 
tial ingredient. 

Alloys  of  Tin  containing  a  minor  percentage  of  tin  are 
bronze,  gun  metal,  speculum  metal,  type  metal,  anti-fric- 
tion metals,  and  fusible  alloys.  Britannia  metal  contains 
about  90  per  cent  tin,  8  per  cent  antimony,  and  the  rest 
mainly  copper.  It  is  a  white  metal,  and  was  formerly  made 
into  tableware.  White  metal  contains  less  tin  and  more 
antimony  than  Britannia,  though  the  composition  varies. 
It  resembles  Britannia.  The  harder  varieties  of  white  metal 
are  used  as  parts  of  machinery,  and  the  softer  kinds  are  made 
into  ornaments  and  cheap  jewelry.  Pewter  and  solder 
contain  varying  proportions  of  tin  and  lead.  Plumbers' 
solder,  or  soft  solder,  is  about  one  third  tin  and  two  thirds 
lead;  it  is  harder  than  either  constituent,  but  melts  at  a 
lower  temperature.  Tin  amalgam  is  sometimes  used  to 
coat  mirrors. 

Compounds  of  Tin.  —  Tin  forms  two  series  of  compounds, 
the  stannous  and  the  stannic.  The  valence  of  tin  is  two  in 
the  stannous  compounds  and  four  in  the  stannic.  Stannic 
oxide  (Sn02)  has  already  been  mentioned  as  the  chief  ore 
of  tin  and  the  product  formed  when  tin  is  burned.  The 
artificial  oxide  is  faint  yellow  when  hot  and  white  when  cold. 
The  native  oxide  is  a  brown  or  black,  lustrous,  (often)  crys- 
talline solid.  Irregular  pebbles  called  stream  tin  occur 
in  some  localities  near  rivers.  Stannous  chloride  (SnCl2) 
is  formed  by  the  interaction  of  hydrochloric  acid  and  tin. 
From  the  concentrated  solution  a  greenish  salt  crystallizes 
(SnCl2. 2  H2O) ,  known  as  the  tin  crystals  or  salt  of  tin. 


TIN  509 

Stannous  chloride  passes  readily  into  stannic  chloride 
(SnCl4)  when  added  to  mercuric  chloride  solution.  The 
simplest  equation  for  this  change  is  — 

SnCl2     +     2HgCl2     =     SnCl4     +       2  HgCl 

•  Stannous  Mercuric  Stannic  Mercurous 

Chloride  Chloride  Chloride  Chloride 

By  an  extension  of  the  simplest  idea  of  oxidation  and  reduc- 
tion, the  stannous  chloride  in  this  change  is  said  to  be 
oxidized  to  stannic  chloride  and  the  mercuric  chloride  to  be 
reduced  to  mercurous  chloride.  An  excess  of  stannous 
chloride  reduces  the  white  mercurous  chloride  to  a  gray 
precipitate  of  finely  divided  mercury;  this  reaction  is  used 
as  a  test  for  tin.  (Compare  tests  for  Mercury,  Chapter 
XXVIII.)  Stannous  chloride  is  often  used  as  a  reducing 
agent  and  as  a  mordant  in  calico  dyeing  and  printing.  Stan- 
nic chloride  (SnCl4)  is  a  colorless,  fuming  liquid;  it  forms 
a  crystalline  hydrate  (SnCl4. 5H20),  known  commercially 
as  oxymuriate  of  tin,  which  is  used  as  a  mordant.  Am- 
monium stannic  chloride  ((NH^SnCle),  or  "pink  salt/' 
is  also  used  as  a  mordant.  Tin  mordants  produce  brilliant 
colors.  Sodium  stannate  (Na2SnO3 . 3  H2O)  is  extensively 
used  to  prepare  cotton  cloth  for  printing.  With  hydrogen 
sulphide,  stannous  compounds  form  brown  stannous  sulphide 
(SnS),  and  stannic  compounds  form  yellow  stannic  sulphide 
(SnS2) ;  both  sulphides  dissolve  in  ammonium  polysulphide, 
owing  to  the  formation  of  soluble  sulpho-salts  of  tin. 

Miscellaneous.  —  Tin,  like  zinc  and  aluminium,  acts  both 
as  a  metal  and  a  non-metal.  Thus,  there  are  the  stannous 
and  stannic  salts  and  the  stannates.  Solutions  of  stannous 
salts  contain  stannous  ions  (Sn++). 

The  atomic  weight  of  tin  is  119.0,  and  the  valence,  as 
stated  above,  is  two  in  stannous  and  four  in  stannic  com- 
pounds. 


510  INORGANIC  CHEMISTRY 

LEAD 

Occurrence  of  Lead.  —  Metallic  lead  is  occasionally  found 
in  small  quantities.  The  most  abundant  ore  is  lead  sul- 
phide (galena,  PbS).  Other  native  compounds  are  the  car- 
bonate (cerussite,  PbCO3),  the  sulphate  (anglesite,  PbS04), 
and  the  phosphate  (pyromorphite,  Pb5Cl(PO4)3). 

Lead  has  been  used  by  civilized  people  since  the  dawn  of  history. 
The  Chinese  have  used  it  for  ages  to  line  chests  in  which  tea  is  stored 
and  transported.  The  Romans  called  it  plumbum  nigrum,  i.e.  black 
lead.  The  symbol  Pb  comes  from  plumbum. 

Lead  ores  are  found  in  the  United  States  mainly  in  the  Middle 
West,  Colorado,  Idaho,  and  Utah. 

Metallurgy  of  Lead.  —  Lead  is  obtained  from  galena  by 
several  processes.  (1)  Ores  rich  in  lead  are  roasted  in  a 
reverberatory  furnace  until  a  part  of  the  sulphide  is  changed 
into  lead  oxide  and  lead  sulphate.  Thus  :  — 


2  PbS     - 

Lead 
Sulphide 

f       302       = 
Oxygen 

2PbO     - 

Lead 
Oxide 

h     2S02 

Sulphur 
Dioxide 

PbS 

Lead 
Sulphide 

f      202     = 

Oxygen 

PbSO4 

Lead 
Sulphate 

The  air  is  then  excluded  and  the  temperature  raised ;    the 
mixture  interacts  thus  :  — 

2  PbS     +      PbSO4     +     2PbO     =      5Pb     +     3SO2 

Lead  Lead  Lead  Lead  Sulphur 

Sulphide  Sulphate  Oxide  Dioxide 

(2)  Ores  poor  in  lead  are  sometimes  reduced  by  roasting 
with  iron,  the  equation  for  the  reaction  being  :  — 

PbS     +     Fe     =      Pb     +     FeS 

Lead  Iron  Lead  Iron 

Sulphide  Sulphide 

(3)  Lead  ores  rich  in  silver  are  roasted  and  then  reduced  with 
coal  (or  coke),  limestone,  and  iron  ore. 


LEAD  511 

Lead  is  refined  by  heating  it  to  oxidize  most  of  the  copper, 
arsenic,  and  antimony,  and  then  treating  the  alloy  by  the 
Parkes  process  (page  453).  In  an  electrolytic  process  the 
cathode  is  a  sheet  of  pure  lead,  the  anode  is  a  plate  of  impure 
lead,  and  the  electrolytic  solution  is  a  mixture  of  lead  fluo- 
silicate  (PbS  iF6)  and  gelatin ;  pure  lead  is  deposited  on  the 
cathode  and  most  of  the  other  metals  remain  attached  to 
the  remna  t  of  the  anode,  from  which  they  are  subsequently 
recovered,  especially  the  gold  and  silver. 

Properties  of  Lead.  —  Lead  is  a  blue-gray  metal.  When 
scraped  or  cut,  it  has  a  brilliant  luster,  which  soon  dis- 
appears, owing  to  the  formation  of  a  film  of  oxide  or  other 
lead  compound.  This  coating  protects  the  lead  from  further 
change.  It  is  a  soft  metal,  and  can  be  scratched  with  the 
finger  nail.  It  discolors  the  hands,  and  when  drawn  across 
a  rough  surface  it  leaves  a  black  mark.  For  this  reason  it 
is  sometimes  erroneously  called  black  lead.  (See  Graphite.) 
Lead  is  not  tough  nor  very  ductile,  though  it  can  be  made  into 
wire,  pressed  (while  soft)  into  pipe,  and  rolled  into  sheets. 
It  is  a  heavy  metal,  its  specific  gravity  being  about  11.4; 
with  the  exception  of  mercury,  it  is  the  heaviest  of  the  familiar 
metals.  It  melts  at  327°  C.,  or  about  100°  higher  than  tin 
and  100°  lower  than  zinc.  Lead,  when  heated  strongly 
in  the  air,  changes  into  lead  monoxide  (litharge,  PbO) ;  at 
higher  temperatures  (about  350°  C.  and  above)  the  tetroxide 
(PbaO4)  is  formed.  Hydrochloric  and  sulphuric  acids  exert 
only  a  slight  chemical  action  upon  compact  lead.  (See  Lead 
Sulphate.)  Nitric  acid  changes  it  into  lead  nitrate  (Pb(NO3)2). 
Acetic  acid  (or  vinegar)  and  acids  from  fruits  and  vegetables 
change  it  into  soluble,  poisonous  compounds ;  hence  cheap 
tin-plated  vessels,  which  sometimes  contain  lead,  should 
never  be  used  in  cooking.  Zinc  and  iron  precipitate  lead  from 
its  solutions  as  a  grayish  mass,  which  often  has  a  beautiful 
treelike  appearance. 


512  INORGANIC  CHEMISTRY 

This  displacement,  as  in  the  case  of  tin,  is  due  to  the  fact 
that  lead  is  lower  than  zinc  and  iron  in  the  electromotive 
series. 

Lead  in  Drinking  Water.  —  Lead  is  slowly  changed  into 
soluble  compounds  by  water  containing  free  oxygen,  carbon 
dioxide,  ammonia,  nitrates,  or  chlorides.  But  water  con- 
taining sulphates  or  carbonates  forms  an  insoluble  coating 
on  the  lead,  thus  protecting  it  from  further  action.  All 
lead  salts  are  poisonous;  and  if  taken  into  the  system  they 
slowly  accumulate  and  ultimately  cause  serious  and  danger- 
ous illness.  Water  suspected  of  attacking  lead  should 
never  be  drunk  after  it.  has  been  standing  very  long  in  lead 
pipes,  but  should  be  allowed  to  flow  until  the  pipe  has  been 
filled  with  fresh  water.  It  is  sometimes  safer  to  substitute 
an  iron  or  block  tin  pipe  for  the  customary  lead  service  pipe. 

Uses  of  Lead.  —  Lead  on  account  of  its  plasticity  when 
warm  is  extensively  made  into  pipe,  which  can  be  easily 
cut,  bent,  and  united  (by  solder).  Lead  pipe  is  not  only 
used  to  convey  water  to  and  from  parts  of  buildings,  but 
as  a  sheath  for  electric  wires,  both  overhead  and  under- 
ground. In  the  form  of  sheets  it  is  used  to  cover  roofs 
and  to  line  sinks,  cisterns,  and  the  cells  employed  in  some 
electrolytic  processes.  The  lead  chambers  and  evaporating 
pans  used  in  manufacturing  sulphuric  acid  are  made  of 
sheet  lead.  Shot  and  bullets  are  lead  (alloyed  with  a  little 
arsenic) . 

The  Alloys  of  Lead  are  important.  Type  metal  contains 
70  to  80  per  cent  lead;  the  other  constituents  are  tin  and 
antimony.  The  alloy  is  harder  than  the  lead  itself  and 
expands  on  cooling,  thereby  making  the  face  of  the  type 
sharp  and  hard.  Solder,  pewter,  and  fusible  alloys  contain 
lead  as  an  essential  constituent.  (See  these  Alloys.) 


LEAD  513 

Lead  Oxides.  —  There  are  three  important  oxides.  Lead 
monoxide  (PbO)  is  a  yellowish  powder  known  as  massicot, 
or  a  buff-colored  crystalline  mass  called  litharge.  It  is 
formed  by  heating  lead  above  its  melting  point  in  a  current 
of  air.  It  is  made  this  way,  though  considerable  is  obtained 
as  a  by-product  in  separating  silver  from  lead.  (See  Cupella- 
tion.)  Large  quantities  are  used  in  preparing  certain  oils 
and  varnishes,  flint  glass,  other  lead  compounds,  and  as  a 
glaze  for  pottery.  Lead  tetroxide  (red  lead  or  minium, 
Pb3O4)  is  a  red  powder,  varying  somewhat  in  color  and 
composition.  It  is  prepared  by  heating  lead  (or  lead  mo- 
noxide) to  about  350°  C.  It  is  used  in  making  flint  glass. 
Pure  grades  are  made  into  artists'  paint,  but  the  cheap 
variety  is  used  to  paint  structural  iron  work  (bridges,  gas- 
ometers, etc.),  hulls  of  vessels,  and  agricultural  implements. 
A  mixture  of  red  lead  and  oil  is  used  in  plumbing  and  gas 
fitting  to  make  joints  tight.  Orange  mineral  has  the  same 
composition  as  red  lead,  although  its  color  is  lighter;  its 
uses  are  the  same.  Lead  dioxide  (lead  peroxide,  PbO2)  is 
a  brown  powder  formed  by  treating  lead  tetroxide  with 
nitric  acid  or  by  the  action  of  chlorine  on  an  alkaline  solution 
of  lead  acetate.  It  is  used  in  storage  batteries. 

Lead  Carbonate,  PbC03,  is  found  native  as  the  trans- 
parent, crystalline  mineral  cerussite.  It  is  obtained  as  a 
white  powder  by  adding  ammonium  carbonate  solution  to 
lead  nitrate  solution.  Sodium  or  potassium  carbonate, 
however,  forms  basic  lead  carbonates,  whose  composition 
depends  upon  the  temperature.  The  most  important  of 
these  basic  carbonates  usually  has  the  composition  corre- 
sponding to  the  formula  Pb3(OH)2(CO3)2,  and  is  known  as 
white  lead.  It  is  a  heavy  white  powder  which  mixes  well 
with  linseed  oil,  and  is  used  extensively  as  a  white  paint 
and  as  the  basis  of  many  colored  paints.  Its  value  as  a 


514  INORGANIC  CHEMISTRY 

paint  is  largely  due  to  its  superior  covering  power,  i.e.  a 
very  thin  layer  produces  a  perfectly  white,  opaque  surface. 
In  recent  years  other  substances  have  been  mixed  with  or 
substituted  for  white  lead,  e.g.  zinc  oxide,  barium  sulphate, 
and  lithophone.  These  solids  are  white,  do  not  darken  in 
air  (as  white  lead  does),  and  often  improve  the  paint  in  other 
ways. 

White  lead  is  manufactured  by  several  processes.  The  Dutch 
process  is  the  oldest,  having  been  used  as  early  as  1622.  It  is  essen- 
tially the  same  to-day,  though  many  details  have  been  improved. 
Perforated  disks  of  lead  are  put  in  earthenware  pots  which  have  a 
separate  compartment  at  the  bottom  containing  a  weak  solution  of 
acetic  acid  (about  as  strong  as  vinegar).  These  pots  are  arranged 
in  tiers  in  a  large  building,  and  spent  tan  bark  is  placed  between  each 
tier.  The  building  is  now  closed,  except  openings  for  the  entrance 
and  exit  of  air  and  steam.  The  heat  volatilizes  the  acetic  acid, 
which  changes  the  lead  into  a  lead  acetate.  The  tan  bark  ferments 
and  liberates  carbon  dioxide,  which  changes  the  lead  acetate  into 
basic  lead  carbonate.  The  whole  operation  requires  from  sixty  to  one 
hundred  days.  Other  processes,  much  the  same  as  the  Dutch  process, 
are  used,  their  chief  aim  being  to  lessen  the  time  of  manufacture. 

Lead  Sulphide,  PbS.  —  Native  lead  sulphide  is  the  mineral 
galena,  the  chief  ore  of  lead.  It  resembles  lead  in  appear- 
ance, but  is  harder  and  is  usually  crystallized  as  cubes, 
octahedrons,  or  their  combinations.  It  has  a  perfect  cubic 
cleavage,  i.e.  it  breaks  into  cubes  or  fragments  more  or  less 
rectangular.  It  is  easily  changed  into  lead  by  heating  it 
alone  or  with  sodium  carbonate  on  charcoal.  Lead  sulphide 
as  prepared  in  the  laboratory  is  a  black  solid.  Black  lead 
sulphide  is  readily  precipitated  from  a  lead  salt  solution 
by  hydrogen  sulphide.  Its  formation  is  a  test  for  lead.  It 
is  changed  into  lead  chloride  by  concentrated  hydrochloric 
acid  and  into  lead  sulphate  by  concentrated  nitric  acid. 
Additional  tests  for  lead  are  the  formation  of  the  sulphate 
and  chromate,  as  described  in  the  next  paragraph. 


CERIUM   AND  THORIUM  515 

Other  Compounds  of  Lead,  which  are  important,  are  the 
chloride,  sulphate,  nitrate,  chromate,  and  acetate.  Lead 
chloride  (PbCl2)  is  a  white  solid  formed  by  adding  hydro- 
chloric acid  or  a  soluble  chloride  to  a  cold  solution  of  a  lead 
salt.  It  dissolves  in  hot  water.  Lead  sulphate  (PbS04)  is  a 
white  solid  formed  by  adding  sulphuric  acid  or  a  soluble 
sulphate  to  a  solution  of  a  lead  salt.  It  is  very  slightly 
soluble  in  water,  but  soluble  in  sulphuric  acid,  hence  crude 
sulphuric  acid  often  contains  lead  sulphate.  Lead  nitrate 
(Pb(NO3)2)  is  a  white  crystalline  solid  formed  by  dissolving 
lead  (or  better,  lead  monoxide)  in  nitric  acid.  When  heated, 
it  decomposes  into  lead  oxide  (PbO),  nitrogen  peroxide, 
and  oxygen.  Lead  chromate  (PbCrO4)  is  a  yellow  solid 
formed  by  adding  a  solution  of  a  lead  compound  to  a 
solution  of  potassium  chromate  or  potassium  dichromate. 
It  is  sometimes  called  "chrome  yellow."  Lead  acetate 
(Pb(C2H3O2)2)  is  a  white  crystalline  solid  formed  by  the 
action  of  acetic  acid  upon  lead  or  lead  oxide  (PbO).  It 
is  very  soluble  in  water  and  is  often  called  "  sugar  of  lead." 

Miscellaneous.  —  Aqueous  solutions  of  lead  compounds 
may  contain  several  kinds  of  ions.  The  common  com- 
pounds, such  as  lead  chloride,  lead  nitrate,  and  lead  acetate, 
yield  a  colorless  ion  (Pb++).  Other  compounds  yield  ions 
containing  oxygen  as  well  as  complex  ions.  Compounds  in 
which  lead  acts  as  a  non-metal  are  known,  e.g.  sodium 
plumbate  (Na2PbO3). 

The  atomic  weight  of  lead  is  207.10.  The  valence  is  two 
in  many  of  its  compounds;  it  is  four  in  lead  dioxide  (PbO2) 
and  plumbates. 

CERIUM  AND  THORIUM 

Cerium  (Ce)  and  Thorium  (Th)  are  members  of  a  family 
of  rare  elements  in  the  same  periodic  group  as  tin  and  lead. 
They  are  constituents  of  rare  minerals,  but  their  compounds 


516  INORGANIC  CHEMISTRY 

are  prepared  from  a  complex  mineral  substance  named 
monazite  sand.  The  oxides  of  thorium  and  cerium  are  the 
essential  compounds  in  the  Welsbach  mantles.  In  making 
the  mantles,  the  cotton  bag  is  dipped  into  a  solution  of 
thorium  and  cerium  nitrates,  and  then  burned.  The  cotton 
is  destroyed  by  heat  and  the  nitrates  are  converted  into  a 
firm  mass  of  oxides,  which  retain  the  shape  of  the  mantle. 
The  proportion  is  about  1  per  cent  cerium  oxide  (Ce02) 
and  99  per  cent  thorium  oxide  (Th02),  this  being  the  proper 
mixture  for  a  brilliant  flame  of  suitable  color.  (See  also 
Bunsen  Flame,  Chapter  XVI.) 
Thorium  compounds  are  radioactive.  (See  Radioactivity.) 

PROBLEMS  AND   EXERCISES 

1.  What  is  the  per  cent  of  lead  in  (a)  galena  (PbS),  (6)  cerus- 
site  (PbCOa),    (c)    anglesite    (PbSO4),    (d)    crystallized  lead  acetate 
(Pb(C2H3O2)2.3H2O)? 

2.  How  much  litharge  may  be  made  from  40.5  gm.  of  lead? 
(Assume  Pb  +  O  =  PbO.) 

3.  What  is  the  per  cent  of  tin  in  (a)  tinstone  (SnO2),  (6)  stannous 
chloride  (SnCl2),  (c)  stannic  chloride  (SnCU)  ? 

4.  If   100  gm.  of  tin  form  127.1  gm.  of  stannic  oxide   (SnO2), 
what  is  the  atomic  weight  of  tin? 

5.  By  analysis  100  gm.  of  lead  monoxide  yielded  7.1724  gm.  of  oxy- 
gen.    What  is  the  atomic  weight  of  lead  ? 

6.  The  formulas  of  cerium  oxide  and  thorium  oxide  are  CeO2  and 
ThO2 ;  what  are  the  formulas  of  the  corresponding  chlorides,  nitrates, 
and  sulphides? 

7.  Calculate  the  solubility  product  of  lead  chromate  on  the  as- 
sumption that  the  ionization  is  100  per  cent  and  the  molar  solubility 
is  .0000004. 


CHAPTER  XXXI 
Manganese 

Occurrence.  —  This  metal  is  not  found  free  in  nature, 
but  its  oxides  are  widely  distributed  and  rather  abundant. 
The  chief  compound  is  manganese  dioxide  (pyrolusite, 
MnO2).  Other  native  compounds  of  manganese  are  braunite 
(Mn2O3),  hausmannite  (Mn3O4),  manganite  (MnO(OH)  or 
MnO2H),  and  rhodocroisite  (MnC03). 

Preparation,  Properties,  and  Uses.  —  Manganese  is  pre- 
pared by  heating  manganese  dioxide  with  charcoal  in  an 
electric  furnace,  but  the  purest  quality  is  made  by  the 
aluminothermic  method. 

The  equation  for  the  latter  change  is  — 

4A1      +     3MnO2     =      3  Mn     +     2  A12O8 

Aluminium          Manganese  Manganese          Aluminium 

Dioxide  Oxide 

The  metal  is  gray-red,  hard,  and  brittle.  If  pure,  it  is 
permanent  in  the  air.  It  interacts  with  dilute  acids.  Its 
chief  property  is  the  ready  formation  of  alloys,  especially 
with  iron,  copper,  zinc,  and  nickel.  It  melts  at  1225°  C. 

Alloys  of  Manganese  and  iron  are  extensively  used  in  the 
manufacture  of  Bessemer  steel.  (See  Steel.)  Spiegel  iron 
contains  from  5  to  20  per  cent  of  manganese,  while  ferro- 
manganese  contains  20  per  cent  or  more.  Other  alloys  are 
finding  applications,  now  that  pure  manganese  is  available. 

Manganese  Dioxide,  MnO2,  is  the  most  abundant  and  im- 
portant compound.  It  is  a  black  solid  and  is  often  called 

517 


518  INORGANIC  CHEMISTRY 

black  oxide  of  manganese.  When  heated  to  a  high  tem- 
perature, it  yields  oxygen;  and  when  heated  with  hydro- 
chloric acid,  the  two  compounds  interact,  forming  man- 
ganous  chloride,  chlorine,  and  water,  thus :  — 

MnO2    +      4HC1      =     MnCl2     +     C12    +    2H20 

Manganese       Hydrochloric        Manganous          Chlorine        Water 
Dioxide  Acid  ^Chloride 

It  colors  glass  and  borax  a  beautiful  amethyst,  and  is  often 
used  in  glass  making  to  neutralize  the  green  color  that 
would  be  caused  by  iron  compounds  in  the  sand.  Large 
quantities  are  used  in  the  manufacture  of  oxygen,  chlorine, 
glass,  and  manganese  alloys  and  compounds. 

A  borax  bead  is  colored  amethyst  by  manganese  com- 
pounds in  the  oxidizing  flame  but  is  colorless  in  the  reduc- 
ing flame  —  a  test  for  manganese. 

The  manganese  dioxide  used  in  the  manufacture  of  chlorine  is 
recovered  by  the  Weldon  process.  The  impure  manganous  chloride 
solution  from  the  chlorine  still  is  treated  with  calcium  carbonate  to 
neutralize  the  free  acid  and  precipitate  any  iron  present.  Lime  is 
added  to  the  clear  solution  of  manganous  chloride,  and  air  is  blown 
into  the  mixture.  The  manganous  chloride  is  changed  into  man- 
ganous hydroxide  (Mn(OH)2),  which  interacts  with  the  oxygen  (of 
the  air)  and  lime,  forming  chiefly  calcium  manganite  (CaMnO3). 
After  this  mixture  has  settled,  the  calcium  chloride  is  drawn  off, 
and  the  manganese  compound,  which  is  called  "  Weldon  mud,"  is 
used  to  generate  more  chlorine. 

Manganese  dioxide  was  used  by  the  ancients  to  decolorize  glass, 
but  its  nature  was  misunderstood.  They  confused  it  with  an  iron 
oxide  called  magnesia  stone,  and  the  alchemists  in  the  Middle  Ages 
gave  the  name  magnesia  to  this  manganese  dioxide.  Later  they 
called  it  magnesia  nigra,  or  black  magnesia,  to  distinguish  it  from 
magnesia  alba,  or  white  magnesia  (MgO),  supposing  the  two  were 
related.  Manganese  was  isolated  in  1774,  and  later  was  given  the 
specific  name  manganesium,  which  was  soon  shortened  to  manganese. 

Potassium  Permanganate,  KMnO4,  is  a  dark  purple, 
glistening,  crystalline  solid,  though  the  crystals  sometimes 


MANGANESE  519 

appear  black,  with  greenish  luster.  It  is  very  soluble  in 
water,  and  solution  is  red,  purple,  or  black,  according  to 
the  concentration.  Potassium  permanganate  gives  up  its 
oxygen  readily  and  is  frequently  used  as  an  oxidizing  agent 
in  the  laboratory.  It  is  also  used  as  a  disinfectant,  a  medi- 
cine, in  bleaching  and  dyeing,  in  coloring  wood  brown,  and 
in  purifying  gases,  such  as  hydrogen,  ammonia,  and  carbon 
dioxide. 

Potassium  permanganate  is  manufactured  by  oxidizing  a 
mixture  of  manganese  dioxide  and  potassium  hydroxide, 
and  treating  the  resulting  potassium  manganate  with  sul- 
phuric acid,  carbon  dioxide,  or  chlorine.  The  essential  re- 
actions are  represented  thus  :  — 

MnO2    +  2KOH  +  O  =     K2MnO4    +     H2O 

Manganese        Potassium  Potassium 

Dioxide          Hydroxide  Manganate 

3K2MnO4  +  2CO2     =    2KMn04  +  2K2CO3  +  MnO2 

Potassium 
Permanganate 

The  uses  of  potassium  permanganate  depend  mainly  upon 
its  oxidizing  power,  that  is,  upon  the  property  of  liberating 
nascent  oxygen  readily.  In  an  acid  solution  the  action  is 
represented  thus  :  — 

2KMnO4  +  3H2SO4  =  5O  +  2MnSO4  +  KsSO.-f  3H2O 

Potassium  Sulphuric        Oxygen      Manganese  Potas-         Water 

Permanganate  Acid  Sulphate          sium  Sul- 

phate 

In  neutral  or  alkaline  solutions  the  action  is  as  follows :  — 
2  KMn04  +  5  H2O  =  3  O  +  2  Mn(OH)4  +  2  KOH 

The  liberated  oxygen  oxidizes  organic  matter  or  any  other 
oxidizable  substance,  and  the  solution  becomes  brown  or 
colorless,  owing  to  the  reduction  of  the  potassium  per- 
manganate and  the  transformation  into  manganese  com- 
pounds having  a  faint  color  or  none. 


520  INORGANIC  CHEMISTRY 

Compounds  of  Manganese  are  numerous,  often  complex, 
and  closely  related.  There  are  four  oxides  besides  man- 
ganese dioxide.  Three  manganous  compounds  are  impor- 
tant, the  chloride  (MnCl2),  the  sulphate  (MnS04),  and  the 
sulphide  (MnS);  in  these  and  other  manganous  compounds 
the  manganese  plays  the  role  of  a  metal.  The  chloride  and 
sulphate  are  pink,  crystalline  salts.  The  sulphide  is  obtained 
as  a  flesh-colored  precipitate  by  adding  ammonium  sulphide 
to  the  solution  of  a  manganous  salt,  the  color  distinguish- 
ing it  from  all  other  sulphides;  its  formation  is  often  used 
as  a  test  for  manganese.  Manganates  are  salts  of  the  hypo- 
thetical manganic  acid  (H2MnO4);  the  manganese  in  them 
acts  as  a  non-metal.  Potassium  manganate  (K2MnO4)  is 
obtained  as  a  green  mass  by  fusing  a  mixture  of  a  man- 
ganese compound,  potassium  hydroxide  (or  carbonate),  and 
potassium  nitrate.  Its  formation  on  a  small  scale  consti- 
tutes a  test  for  manganese.  One  equation  for  this  chemical 
change  is  given  on  the  preceding  page ;  another  is  as  follows : 

Mn02       +       K2CO3    +    0    =    K2MnO4       +       CO2 

Manganese  Potassium  Potassium  Carbon 

Dioxide  Carbonate  Manganate  Dioxide 

Sodium  manganate  (Na2MnO4)  is  used  in  solution  as  a  disin- 
fectant, but  sodium  permanganate  (NaMnO4)  is  more  effective 
and  is  sold  in  solution  as  "  Condy's  liquid." 

Miscellaneous.  —  Manganese  has  a  variable  valence.  It  is 
two  in  manganous  compounds  (e.g.  MnO,  Mn(OH)2,  MnS, 
and  MnSO4),  three  in  manganic  compounds  (e.g.  Mn203), 
four  in  manganese  dioxide  (MnO2)  and  in  manganites,  six 
in  manganates  (e.g.  K2MnO4),  and  seven  in  permanganates 
(e.g.  KMn04).  The  valence  of  the  radical  MnO4  is  two  in 
manganates  and  one  in  permanganates. 

Manganese  compounds  yield  several  kinds  of  ions,  e.g.  the 


MANGANESE  521 

delicate  pink  manganese  ion  (Mn++),  and   the   purple  per- 
manganate ion  (Mn04~). 

The  atomic  weight  of  manganese  is  54.93.  Manganese 
apparently  occupies  an  isolated  position  in  Group  VII  of  the 
periodic  classification. 

PROBLEMS  AND  EXERCISES 

1.  Calculate  the  weight  of  manganese  in  (a)  1  metric  ton  of  pyro- 
lusite  (85  per  cent  pure),  (6)  1  kg.  of  manganous  chloride,  and  (c)  275 
gm.  of  potassium  manganate. 

2.  Write  the  formulas  of  the  following :    Manganese  carbonate, 
manganese  heptoxide,  barium  permanganate,  manganic  acid,  potas- 
sium manganese  alum.     Calculate  the  per  cent  of  manganese  in 
three  of  these  compounds. 

3.  How  much  manganese  dioxide  is  needed  to  prepare  a  ton 
(2000  Ib.)  of  potassium  permanganate? 

4.  What   (a)  weight  and   (6)  volume   (standard  conditions)  of 
oxygen  are  produced  by  the  interaction  of  sulphuric  acid  and  90  gm. 
of  potassium  permanganate? 

5.  What  is  the  atomic  weight  of  manganese,  if  10.6647  gm.  of 
manganous  oxide  (MnO)  yield  22.6875  gm.  of  manganese  sulphate 
(MnSOO  ?     (Use  exact  atomic  weights.) 

6.  From  the  following  data,  find  the  atomic  weight  of  manganese 
and  the  number  of  atoms  in  the  compound  analyzed :    An  analysis 
of  an  oxide  of  manganese  yielded  69.62  parts  of  manganese  and  30.38 
parts  of  oxygen ;    specific  heat  of  manganese  is  .1217. 


CHAPTER  XXXII 
Chromium  —  Uranium  —  Radioactivity 

Occurrence.  —  Metallic  chromium  is  never  found  free.  Its 
chief  ore  is  ferrous  chromite,  or  chrome  iron  ore,  Fe(CrO2)2. 
Native  lead  chromate  (crocoite,  or  crocoisite,  PbCrO4)  is  less 
common.  Traces  of  chromium  occur  in  many  green  miner- 
als and  rocks,  e.g.  emerald,  serpentine,  and  verde  antique 
marble. 

The  name  chromium  comes  from  the  Greek  chroma,  meaning  color, 
and  emphasizes  the  fact  that  most  chromium  compounds  have  con- 
spicuous colors. 

Preparation,  Properties,  and  Uses.  —  Chromium  was  a 
rare  metal  until  Moissan  prepared  it,  in  1894,  by  heating  a 
mixture  of  chromite  and  carbon  in  an  electric  furnace.  The 
product  contained  carbon  and  was  refined  by  fusing  it  with 
lime.  Pure  chromium  is  now  prepared  by  the  alumino- 
thermic  method.  (See  Thermit.) 

Chromium  is  a  lustrous  gray  metal.  It  takes  a  good 
polish,  which  is  not  removed  by  exposure  to  air.  It  is  hard 
and  brittle,  and  can  be  polished  without  difficulty.  Its 
specific  gravity  is  about  6.9.  Its  melting  point  is  1510°  C. 
The  heated  metal  burns  in  oxygen,  forming  green  chromic 
oxide  (Cr2O3). 

Chromium  is  used  to  harden  the  steel  that  is  made  into 
armor  plates,  hard  tools,  projectiles,  safes  and  vaults,  and 
certain  parts  of  machines  used  to  crush  gold-bearing  quartz. 
This  hardened  steel  is  called  chromium  steel.  The  com- 
mercial form  of  chromium  is  usually  an  alloy  of  65  to  80 

522 


CHROMIUM  523 

per  cent  chromium,  a  little  carbon,  and  the  rest  iron;  this 
alloy  is  called  ferrochrome. 

Compounds  of  Chromium  are  numerous,  some  are  com- 
plex, many  pass  readily  into  one  another,  and  a  few  have 
industrial  applications.  The  most  important  are  potassium 
chromate,  potassium  dichromate,  chrome  alum,  and  lead 
chromate. 

Potassium  Chromate,  K2CrO4,  and  Potassium  Dichromate, 
or  Bichromate,  K2Cr2O7.  —  These  compounds  are  manu- 
factured from  chrome  iron  ore.  The  crushed  ore  is  mixed 
with  lime  and  potassium  carbonate  and  roasted  in  a  rever- 
beratory  furnace;  air  is  freely  admitted  and  the  mass  fre- 
quently raked.  By  this  operation  the  ore  is  oxidized  into  a 
mixture  of  calcium  and  potassium  chromates.  The  mass  is 
cooled,  pulverized,  and  treated  with  a  hot  solution  of  potas- 
sium sulphate,  which  changes  the  calcium  chromate  into 
potassium  chromate,  thus  :  — 

CaCrO4  +   K2S04    =    K2CrO4   +   CaSO4 

Calcium  Potassium        Potassium  Calcium 

Chromate  Sulphate          Chromate  Sulphate 

The  solution  of  potassium  chromate  is  filtered,  concentrated, 
and  changed  by  sulphuric  acid  into  potassium  dichromate; 
the  latter  is  purified  by  recrystallization  from  water.  Potas- 
sium chromate  is  a  lemon-yellow,  crystalline  solid,  which 
contains  no  water  of  crystallization.  It  is  very  soluble  in 
water  and  gives  a  yellow  solution.  Acids  change  it  into  the 
dichromate,  thus  :  — 

2K2CrO4   4-    H2SO4   =  K2Cr2O7   +    K^SO,   +  H2O 

Potassium          Sulphuric         Potassium         Potassium        Water 
Chromate  Acid  Dichromate          Sulphate 

Potassium  chromate  is  also  formed  as  a  yellow  mass  by 
fusing  intensely  on  porcelain  (or  platinum)  a  mixture  of  a 


524  INORGANIC  CHEMISTRY 

chromium  compound,  potassium  carbonate,  and  potassium 
nitrate.  When  the  mass  is  boiled  with  an  excess  of  acetic 
acid  to  decompose  the  carbonate  and  expel  the  carbon 
dioxide,  and  then  added  to  a  lead  salt  solution,  yellow  lead 
chromate  is  precipitated.  This  experiment  is  often  used  as 
a  test  for  chromium.  (See  also  Lead  Chromate.)  Potassium 
dichromate  is  a  red  solid,  which  is  prepared  as  described 
above.  It  is  less  soluble  in  water  than  potassium  chromate, 
and  yields  a  pale  yellow  or  red  solution  according  to  the 
concentration.  Hydroxides  change  it  into  the  chromate, 
thus  :  — 

K2Cr2O7    -f    2KOH  =  2K2CrO4  +  H2O 

Potassium          Potassium          Potassium          Water 
Dichromate         Hydroxide          Chromate 

Potassium  dichromate  is  used  in  dyeing,  calico  printing, 
and  tanning,  in  bleaching  oils,  and  in  manufacturing  chro- 
mium compounds  and  dyestuffs.  Its  uses  depend  mainly 
upon  the  fact  that  it  is  an  oxidizing  agent.  When  potas- 
sium dichromate  and  sulphuric  acid  are  mixed,  the  equation 
for  the  reaction  may  be  written  :  — 

K2Cr2O7  +  4H2S04  =    3O    +  K2SO4  -f  Cr2(SO4)3  +  4  H2O 

Potassium         Sulphuric        Oxygen      Potassium       Chromium 
Dichromate  Acid.  Sulphate          Sulphate 

Some  oxidizable  substance,  however,  such  as  ferrous  sul- 
phate, must  be  present  to  use  up  the  liberated  oxygen, 
thus  :  — 

6FeSO4  -f  3H2SO4  +  3O  =  3Fe2(SO4)3  +  3H2O 

Ferrous  Ferric 

Sulphate  Sulphate 

The  complete  chemical  change  is  often  expressed  as  follows:  — 


K2Cr207  +  7  HaSO,  +  6  FeS04  =  3  Fe2(SO4)3  +  K2SO4 

+  Cr2(S04)3 


CHROMIUM  525 

Potassium  chromate  and  dichromate  are  being  replaced  somewhat 
by  the  corresponding  sodium  salts,  because  the  latter  are  cheaper, 
more  soluble,  and  have  analogous  properties.  The  potassium  salts  are 
anhydrous,  but  crystallized  sodium  chromate  is  Na2CrC>4  .  10  H2O 
and  crystallized  sodium  dichromate  is  Na2Cr2O7  •  2  H2O. 


Chrome  Alum,  KaC^SO^.  24  H2O,  is  a  purple,  crystal- 
line solid.  It  is  analogous  in  composition  and  similar  in 
properties  to  ordinary  alum,  but  it  contains  chromium  in- 
stead of  aluminium.  It  can  be  prepared  by  mixing  potas- 
sium and  chromium  sulphates  in  the  proper  proportion,  or 
by  passing  sulphur  dioxide  into  a  solution  of  potassium 
dichromate  containing  sulphuric  acid.  The  commercial  sub- 
stance is  a  by-product  obtained  in  the  manufacture  of  the 
dyestuff  alizarin.  Chrome  alum  is  used  as  a  mordant  in 
dyeing  and  calico  printing  and  in  tanning. 

Lead  Chromate,  PbCr04,  is  a  bright  yellow  solid,  formed 
by  adding  potassium  chromate  or  dichromate  to  a  solution 
of  a  lead  salt.  An  equation  for  the  chemical  change  is  — 

K2Cr2O7  +  2  Pb(NO3)2  +  H2O  =  2  PbCrO4  +  2  KNO3  -f-  2  HNO3 

Potassium  Lead  Water  Lead  Potassium         Nitric 

Dichromate        Nitrate  Chromate  Nitrate  Acid 

It  is  known  as  chrome  yellow,  and  is  used  in  making  yellow 
paint.  When  boiled  with  sodium  hydroxide,  lead  chromate 
is  changed  into  a  basic  chromate  (PbCrO4  .  PbO  .  H2O)  called 
chrome  red  or  chrome  orange,  depending  on  the  color. 

The  precipitation  of  lead  chromate  by  the  interaction  of 
a  dissolved  lead  salt  and  a  dissolved  chromate  (or  dichro- 
mate) is  often  used  as  a  test  for  chromium. 

Chromium  forms  Four  Series  of  Compounds,  the  chro- 
mous,  the  chromic,  the  chromites,  and  the  chromates 
(mono-  and  di-).  In  the  chromous  and  chromic  compounds, 
chromium  acts  as  a  metal,  but  in  chromites  and  chromates 
it  acts  as  a  non-metal. 


526  INORGANIC  CHEMISTRY 

Chromous  Compounds  may  be  regarded  as  derived  from 
chromous  oxide  (CrO).  As  a  class  they  are  so  easily  oxidized 
into  chromic  compounds  that  they  are  difficult  to  prepare 
and  keep. 

Chromic  Compounds  may  be  regarded  as  derivatives  of 
chromic  oxide  (Cr2O3).  This  is  a  bright  green  powder  pre- 
pared by  heating  chromic  hydroxide  (Cr(OH)3),  and  is  the 
basis  of  the  chrome  green  pigments  used  to  color  glass 
and  ornament  porcelain.  When  chromium  compounds  are 
heated  with  borax,  they  color  the  bead  green  in  both  flames, 
owing  to  the  formation  of  this  oxide  (Cr2O3).  If  potassium 
dichromate  and  boric  acid  are  mixed  and  heated,  and  then 
treated  with  water,  a  hydrated  chromic  oxide  is  formed 
called  Guignet's  green  (Cr2O3 .  2  H2O) ;  it  gives  a  permanent 
color  and  is  extensively  used.  There  are  several  chromic 
hydroxides.  The  typical  one  has  the  composition  repre- 
sented by  the  formula  Cr(OH)3.  It  is  a  bluish  solid  formed 
by  the  interaction  of  a  chromic  compound  (e.g.  chrome 
alum)  and  an  alkaline  hydroxide  or  sulphide.  The  precipi- 
tation of  chromic  hydroxide  by  ammonium  sulphide  is  due 
to  the  fact  that  chromic  sulphide  (which  we  might  expect 
to  be  formed)  hydrolyzes,  yielding  chromic  hydroxide  and 
hydrogen  sulphide.  (Compare  Aluminium  Hydroxide,  Chap- 
ter XXIX.) 

Chromic  hydroxide  is  soluble  in  an  excess  of  sodium 
(or  potassium)  hydroxide.  That  is,  it  is  changed  into 
a  soluble  chromite,  just  as  aluminium  hydroxide  forms 
soluble  aluminates.  Unlike  aluminates,  however,  the  chro- 
mites  are  changed  back  into  chromic  hydroxide  by  boiling. 
Other  chromic  salts  are  chromic  chloride  (CrCl3),  chromic 
sulphate  (Cr2(SO4)3),  and  potassium  chromium  sulphate  or 
chrome  alum  (K2Cr2(SO4)4 .  24  H2O).  The  valence  of  chro- 
mium in  chromic  compounds  is  three, 


URANIUM  527 

Chromites  may  be  regarded  as  salts  of  an  acid  having  the 
composition  corresponding  to  HCrO2;  native  chromite 
(FeO2O4  or  Fe(CrO2)2)  is  an  iron  salt  of  this  acid. 

Chromates  and  dichromates  start  theoretically  from  chro- 
mium tri oxide  (CrO3).  This  is  the  anhydride  of  the  hypo- 
thetical chromic  acid  (H2Cr04).  When  concentrated  sul- 
phuric acid  is  added  to  a  saturated  solution  of  potassium 
dichromate  (or  chromate),  chromium  trioxide  (CrO3)  sepa- 
rates as  long,  bright  red  crystals,  thus:  — 

K20207     +    H2S04      =  2  Cr03     +      K2S04    +     H2O 

Potassium  Sulphuric  Chromium  Potassium  Water 

Dichromate  Acid  Trioxide  Sulphate 

It  is  sometimes  called  chromic  acid,  and  is  a  vigorous 
oxidizing  agent.  The  valence  of  chromium  in  chromates 
and  dichromates  is  six;  the  radicals  CrO4  and  O2O7  have 
the  valence  two. 

Miscellaneous.  —  Chromium  compounds  yield  several  kinds 
of  ions,  e.g.  the  violet  chromic  ion  (Cr+  +  +  ),  the  yellow 
chromate  ion  (Cr04  ),  and  the  red  dichromate  ion 
(CrA— )• 

Molybdenum  (Mo),  Tungsten  (W),  and  Uranium  (U)  are  rare  metal- 
lic elements  related  to  chromium.  Most  of  their  compounds  have 
only  scientific  interest,  though  some  have  analytical  or  industrial  uses. 
Ammonium  molybdate  ((NHOaMoOO  is  used  in  the  laboratory  to  de- 
termine the  amount  of  phosphorus  in  fertilizers  and  iron.  Tungsten 
is  used  to  harden  steel,  and  as  the  filament  of  electric  light  bulbs; 
sodium  tungstate  (NaaWO4)  finds  application  in  rendering  cloth  fire- 
proof. Uranium  compounds  are  obtained  from  uraninite  and  pitch- 
blende, the  latter  mineral  now  being  the  chief  source.  The  element 
forms  many  compounds  and  the  important  ones  are  the  oxides,  sodium 
uranate  (Na2U2O7 .  6  H2O),  and  uranyl  nitrate  (UO2(NO3)2  .  6  H2O); 
from  the  nitrate  other  salts  are  prepared.  Sodium  uranate  is  some- 
times called  uranium  yellow,  and  it  is  used  to  make  uranium  glass. 
Such  glass  is  green  by  transmitted  light  and  yellow  by  reflected  light. 
Uranium  is  a  radioactive  element.. 


528  INORGANIC  CHEMISTRY 


PROBLEMS  AND  EXERCISES 

1.  What  is  the  per  cent  of  chromium  in    (a)  lead  chromate 
(PbCrO4),    (6)    chrome   ironstone    (Fe(CrO2)2),    (c)    chromic   oxide 
(Cr203)? 

2.  How  many  grams  of  lead  chromate  can  be  made  from  438  gm. 
of  potassium  dichromate  ? 

3.  What  is  the  solubility  product  of  barium  chromate  if  the 
ionization  is  89  per  cent  and  the  molar  solubility  is  .000014? 


RADIOACTIVITY 

Historical.  —  It  was  found  about  1896  that  uranium  com- 
pounds affect  a  photographic  plate  and  discharge  an  electro- 
scope. Numerous  experiments  by  Becquerel,  Curie,  Ruther- 
ford, and  others  showed  that  these  effects  were  probably  due 
to  radiations  like  X-rays  and  Rontgen  rays.  Furthermore, 
experiments  indicated  that  this  power  of  radiation  belongs 
to  the  uranium  itself,  and  that  the  radiations  are  emitted 
spontaneously  without  the  aid  of  any  outside  agency.  Other 
substances,  subsequently  shown  to  be  radium  compounds, 
were  found  to  possess  similar  properties,  and,  as  stated  in 
the  preliminary  discussion  of  radium,  such  substances  are 
said  to  exhibit  radioactivity.  Soon  after  (1898)  an  exten- 
sive examination  of  many  minerals  showed  that  certain  min- 
erals containing  uranium,  especially  pitchblende,  are  more 
radioactive  than  uranium  itself.  A  radioactive  substance 
was  extracted  from  pitchblende  by  M.  and  Mme.  Curie,  and 
the  elementary  constituent  in  it  was  named  radium. 

Interpretation  of  Radioactivity.  —  Radioactivity  is  not  due, 
as  was  first  supposed,  to  radiations,  but  to  the  spontaneous 
emission  by  radioactive  substances  of  two  kinds  of  particles, 
which  are  called  alpha  (a)  and  beta  ((3) ;  the  emission  of  beta 
particles  is  accompanied  by  pulsations  in  the  ether  known  as 
gamma  (y)  rays, 


MADAME    CURIE 


RADIOACTIVITY  529 

The  alpha  particles  are  shot  off  in  a  stream  which  moves 
with  a  velocity  averaging  about  18,000  miles  a  second.  The 
path  of  single  alpha  particles  in  a  special  apparatus  has  been 
photographed.  Alpha  particles  bear  a  positive  charge  of 
electricity  and  can  be  readily  detected  by  a  delicate  electro- 
scope. Many  of  the  electrical  phenomena  of  radioactive 
substances  are  due  to  alpha  particles.  Alpha  particles  are 
fdur  times  as  heavy  as  hydrogen  atoms,  and  they  are  iden- 
tical with  charged  atoms  of  the  element  helium. 

An  instrument  called  the  spinthariscope  shows  vividly  that  alpha 
particles  are  being  shot  off  continuously  by  a  radium  compound.  It 
is  a  small  microscope  with  a  screen  and  pin  opposite  the  lens ;  the 
screen  is  coated  with  zinc  sulphide  and  on  the  needle  there  is  a  min- 
ute quantity  of  radium  bromide.  Upon  looking  through  the  lens, 
minute  flashes  of  light  are  seen  on  the  screen.  They  are  due  to  the 
alpha  particles  which  produce  fluorescence  in  the  zinc  sulphide. 

The  beta  particles  consist  of  a  stream  of  electrons,  i.e.  par- 
ticles of  negative  electricity,  moving  with  a  varying  velocity 
which  is  sometimes  nearly  as  great  as  the  velocity  of  light 
(186,000  miles  a  second).  Beta  particles  are  very  light,  their 
weight  being  about  1 810'0  of  the  weight  of  a  hydrogen  atom. 
To  the  beta  particles  are  ascribed  most  of  the  photographic 
effects  of  radioactive  substances. 

The  gamma  rays  are  not  material  particles,  but  like  X-rays 
are  pulsations  in  the  ether.  The  curative  effect  of  radium  is 
believed  to  be  due  to  gamma  rays. 

Alpha  particles  move  in  straight  lines  and  penetrate  air 
to  a  depth  of  3  to  8  centimeters.  They  are  almost  entirely 
stopped  by  a  thin  sheet  of  paper  and  by  aluminium  leaf 
.1  millimeter  thick.  Beta  particles  move  in  straight  lines  at 
first,  but  soon  in  curved  lines,  owing  to  collisions  with  the 
relatively  heavier  molecules  of  the  gases  of  the  air.  They 
penetrate  air  to  a  less  depth  than  alpha  particles  ;  they  pass 
through  gold  leaf  but  are  stopped  by  aluminium  1  centimeter 


530  INORGANIC  CHEMISTRY 

thick.  The  gamma  rays  are  the  most  penetrating.  They 
pass  readily  through  thick  layers  of  metals ;  glass  tubes  con- 
taining radium  salts  are  enclosed  in  lead  vessels  to  absorb 
the  gamma  rays.  The  stoppage  of  so  many  and  such 
rapidly  moving  particles  by  the  air,  metals,  and  the  radio- 
active substance  itself  develops  heat.  It  is  estimated  that 
one  gram  of  radium  (metal,  in  combination)  would  produce 
spontaneously  about  120  calories  per  hour. 

Disintegration  of  Radioactive  Compounds.  —  Uranium  and 
radium  are  like  most  of  the  other  chemical  elements  in  general 
respects,  but  they  differ  in  one  conspicuous  property,  viz. 
atomic  instability.  That  is,  the  atoms  of  uranium,  radium, 
and  a  few  other  elements  are  spontaneously  disintegrating. 
Uranium  is  the  parent,  so  to  speak,  of  a  series  of  products 
formed  step  by  step  by  disintegration.  Among  these  prod- 
ucts are  the  elements  radium  and  niton.  The  rate  of  dis- 
integration varies  widely  and  is  not  affected  by  conditions, 
that  is,  it  is  spontaneous.  In  a  unit  time  a  definite  fraction 
of  a  product  disintegrates ;  it  is  customary  in  describing  dis- 
integration to  state  the  time  in  which  half  the  amount  would 
disintegrate.  Thus,  the  half  period  of  radium  is  2400  years 
and  of  niton  is  5.55  days. 

Certain  interesting  facts  bearing  on  disintegration  are  well 
established.  For  example,  uranium  ores  contain  amounts  of 
radium  that  are  proportional  to  the  uranium.  Again,  a  ura- 
nium compound  gradually  recovers  its  whole  radioactivity 
after  the  radium  has  been  removed.  Furthermore,  helium 
is  produced  at  different  steps  in  the  disintegrating  process ; 
the  helium  given  off  by  radium  compounds  has  been  collected, 
studied,  and  its  rate  of  production  measured  (164  cubic 
millimeters  per  gram  of  radium  (metal)  a  year) . 

The  products  formed  by  successive  disintegration  in  the 
uranium  series  include  electrons  and  helium  besides  radium, 


RADIOACTIVTY  531 

niton,  and  certain  less  known  elements.  In  other  words,  the 
product  of  the  disintegration  of  an  atom  of  an  element  may 
be  an  atom  of  another  element  or  an  atom  plus  either  an  elec- 
tron or  a  charged  helium  atom.  The  uranium  (Uranium 
-  1)  series  is  as  follows  :  — 

Uranium  - 1  (238)  — >-  Uranium  -  Xi  (234)  — *- 
Uranium  -X2  (234)  — >-  Uranium  -2  (234)    — >• 

Ionium  (230)  — >-  Radium  (226)  — >- 

Niton  (222)  — >-  Radium  -A :  (218)      — >• 

Radium-B  (214)  — ^Radium-C  (214)     — >- 

Radium -Ci  (214)  — >-  Radium  -D  (210)     — >- 

Radium -E  (210)  — >-  Radium -F  (210)      — *- 
Radium -G  (206). 

In  the  above  series  it  will  be  noticed  that  the  atomic  weight 
(given  in  the  parenthesis)  of  certain  elements  is  4  less  than 
the  atomic  weight  of  the  preceding  element,  e.g.  niton  (222) 
and  radium  (226).  In  such  cases  the  atomic  disintegration  is 
accompanied  by  the  expulsion  of  a  charged  helium  atom  (i.e. 
an  alpha  particle)  having  the  atomic  weight  4.  On  the  other 
hand,  the  atomic  weight  of  certain  elements  is  the  same,  e.g. 
radium  — D,  —  E,  —  F  (210).  In  these  cases  an  electron  (i.e. 
a  beta  particle)  is  expelled.  Although  certain  elements 
have  the  same  atomic  weight,  they  are  distinct  chemical  ele- 
ments and  form  analogous  compounds  which  have  different 
chemical  properties.  Moreover,  certain  elements  in  this 
series  have  chemical  properties  identical  with  those  of  ele- 
ments already  known ;  it  is  believed  that  radium  —  G  is  the 
final  product  of  disintegration. 

Other  Radioactive  Elements.  —  Thorium  is  radioactive, 
though  to  a  much  less  degree  than  radium.  It  furnishes  a 
series  of  disintegration  products  similar  to  uranium.  Ac- 
tinium and  polonium  are  also  radioactive  elements. 


CHAPTER  XXXIII 

Iron,   Nickel,   and   Cobalt 

IRON 

IRON  is  the  most  useful  of  all  metals.  It  has  been  known 
for  ages,  though  not  so  long  as  the  other  common  metals, 
and  has  been  indispensable  in  the  development  of  the  human 
race. 

The  symbol  of  iron,  Fe,  is  from  the  Latin  word  ferrum.  From 
ferrum  are  derived  the  words  ferric  and  ferrous,  which  give  the 
corresponding  forms  ferri-  and  ferro-  (found  in  such  words  as  ferri- 
cyanide,  ferrocyanide,  etc.). 

Occurrence  of  Iron.  —  Uncombined  iron  is  found  only  in 
meteorites,  which  fall  upon  the  earth  from  remote  regions 
in  space,  and  in  certain  volcanic  rocks.  Combined  iron  is 
abundant  and  widely  distributed,  constituting  about  4.5 
per  cent  of  the  earth's  crust.  It  is  found  in  most  rocks 
and  many  minerals,  in  the  soil,  in  springs  and  natural 
waters,  in  chlorophyll  (the  green  coloring  matter  of  plants), 
and  in  hemoglobin  (the  red  coloring  matter  of  the  blood). 
The  chief  Ores  of  iron  are  hematite  (Fe2O3),  limonite 
(2Fe2O3.3H2O),  magnetite  (FeA),  and  siderite  (FeCO3). 
These  ores  often  contain  some  impurity,  such  as  silica,  clay, 
calcium  or  magnesium  carbonate,  and  small  quantities  of 
compounds  of  sulphur,  phosphorus,  and  manganese. 

Other  abundant  native  compounds  of  iron  are  pyrites 
(FeS2),  pyrrhotite  (varying  from  Fe6S7  to  FenS12),  and  the 
copper-iron  sulphides  (chalcopyrite,  CuFeS2,  and  bornite, 
Cu3FeS3).  They  are  not  used  to  any  extent  as  a  source  of 
iron. 

532 


IRON  533 

The  United  States  leads  the  world  in  the  production  of  iron  ore, 
the  annual  output  for  the  last  few  years  being  about  50,000,000  tons. 
This  vast  quantity  conies  from  twenty-five  different  states,  but  the 
bulk  is  mined  in  Minnesota,  Michigan,  Alabama,  Wisconsin,  Ten- 
nessee, Virginia,  West  Virginia,  and  Colorado.  The  most  abundant 
ore  is  the  red  hematite,  which  comes  chiefly  from  the  Lake  Superior 
region  (Fig.  79);  large  quantities  are  mined  in  Alabama  and  Ten- 
nessee. The  latter  states,  together  with  Virginia  and  West  Virginia, 


from  Greenwich 


FIG.  79.  —  Deposits  of  iron  and  copper  near  Lake  Superior.  The  iron 
regions,  known  as  ranges,  are  Marquette  (1),  Menominee  (2),  Gogebic  (3), 
Vermilion  (5),  Mesabi  (6).  No.  4  is  the  copper  region. 

furnish  most  of  the  limonite  or  brown  iron  ore.  Pennsylvania,  New 
Jersey,  and  New  York  contribute  most  of  the  magnetite,  though 
some  is  mined  also  in  Michigan.  The  carbonate  ores,  which  con- 
stitute less  than  1  per  cent  of  the  output,  come  mainly  from  Ohio, 
Maryland,  and  New  York.  Improvements  in  the  machinery  and 
methods  used  in  mining  and  transporting  iron  ore  have  reduced  its 
cost  and  facilitated  its  production.  Thus,  at  an  incredibly  small 
expense,  ore  from  the  Lake  Superior  region  is  raised  from  open  pits 
by  steam  shovels,  dumped  into  large  cars,  carried  to  shipping  ports 
on  the  lakes,  dumped  again  into  huge  bunkers,  dropped  down  chutes 
into  big  freight  steamers  (many  of  which  hold  6000  tons),  which 
carry  it  to  southern  ports  on  Lake  Michigan,  though  the  large  part 


534 


INORGANIC  CHEMISTRY 


is  sent  to  ports  on  the  south  shore  of  Lake  Erie  and  forwarded  by 
rail  to  Pittsburg,  Pennsylvania.     This  city  is  the  center  of  iron  and 

j  ( steel  industries.    Birmingham, 

Alabama,  is  the  center  of  the 
industry  in  the  South,  because 
near  it  the  necessary  ore,  coal, 
and  limestone  are  conveniently 
located. 

Metallurgy  of  Iron.  — 
Iron  is  extracted  most 
easily  from  its  oxides. 
Therefore  the  ore,  unless 
it  is  very  pure  hematite,  is 
first  crushed  and  roasted  to 
change  it  into  ferric  oxide 
(Fe2O3)  as  far  as  possible. 
The  ore  is  then  mixed 
with  coke  and  limestone 
or  sand,  and  smelted  in  a 
blast  furnace.  The  carbon 
reduces  the  oxide  to  metal- 
lic iron,  which  collects  as  a 
liquid  at  the  bottom  of  the 
furnace  beneath  the  slag 
formed  by  the  limestone 
and  impurities.  The  blast 
furnace  (Fig.  80)  is  a  huge 
circular  tower,  from  forty 
to  ninety  feet  high,  and 
about  thirteen  feet  in  di- 


FIG.  80.  —  Blast  furnace :  A,  throat ;  B, 
bosh  •  C,  crucible  where  the  melted  iron 
collects;  D,  pipes  for  hot-air  blast;  E, 
escape  pipe  for  gases  which  do  not  ameter  at  the  largest  part. 
escape  through  the  "down  comer"; 
G,  cup;  H,  cone;  N,  trough  for  draw- 
ing off  slag;  T,  tuyere;  /,  hole  through 
which  iron  is  withdrawn. 


It  is  built  of  iron  and  lined 
with  fire  brick.    Pipes  near 
the  bottom,  called  tuyeres, 
allow  large  quantities  of  hot  air  to  be  forced  into  the  furnace 


IRON  535 

up  through  the  contents,  thereby  producing  the  high  tem- 
perature required  in  the  smelting;  while  another  pipe  at  the 
top  not  only  permits  the  escape  of  the  hot  gaseous  products, 
but  conducts  them  into  a  series  of  pipes  which  lead  to  dif- 
ferent parts  of  the  plant,  where  the  hot  gases  are  utilized  to 
heat  the  air  which  is  blown  through  the  furnace.  The  blast 
pipes  correspond  to  the  bellows  used  by  the  blacksmith, 
and  the  exit  pipe  to  the  chimney,  except  that  gases  escaping 
through  chimneys  are  usually  wasted. 

When  the  blast  furnace  has  been  heated  to  the  proper 
temperature,  or  is  already  in  operation,  the  mixture  or  charge 
is  carried  to  the  top  by  machinery  and  introduced  into  the 
furnace  by  dumping  it  upon  the  cone-shaped  cover;  the 
weight  lowers  the  cover,  which  flies  back  tightly  into  place 
after  the  materials  roll  into  the  furnace.  The  charge  con- 
sists of  the  proper  mixture  of  ore,  fuel,  and  flux.  The  ore, 
as  stated  above,  is  usually  hematite.  The  fuel  is  coke, 
or  coke  mixed  with  coal.  The  flux  varies  with  the  impurities 
in  the  ore;  thus,  it  is  limestone  if  the  ore  contains  silica  or 
clay,  but  sand  if  the  impurities  are  calcium  or  magnesium 
compounds.  It  is  usually  limestone.  The  object  of  the 
flux  is  to  remove  the  impurities  (just  mentioned)  from  the 
charge  in  the  form  of  readily  fusible  silicates  called  slag. 
As  the  smelting  proceeds,  the  contents  of  the  furnace  slowly 
descend  and  are  changed  into  gases,  iron,  and  slag.  The 
gases  rise  through  the  mass  and  escape  by  pipes,  the  solids 
become  pasty  at  first  and  then  liquid,  the  iron  finally  drop- 
ping through  the  slag  into  the  crucible  at  the  bottom  of  the 
furnace,  where  both  are  tapped  off  through  separate  open- 
ings. Fresh  charges  of  definite  weight  and  proportions 
are  added  at  regular  intervals,  and  the  whole  operation 
continues  without  interruption  for  months  or  even  years. 

The  iron  from  the  furnace  is  usually  poured  into  molds 
of  sand  or  iron  and  allowed  to  solidify.  Such  iron  is  called 


536  INORGANIC  CHEMISTRY 

pig  iron  or  cast  iron.  In  some  plants  the  molten  iron  is 
run  directly  from  the  blast  furnace  into  huge  vessels  called 
converters  and  made  at  once  into  steel  (see  below) ;  in  other 
plants  it  is  kept  molten  in  huge  tanks  until  needed. 

The  chemical  changes  involved  in  the  metallurgy  of  iron 
are  numerous.  In  general,  the  iron  oxide  is  reduced  to 
metallic  iron  largely  by  carbon  monoxide.  The  carbon  of 
the  fuel  at  first  forms  carbon  dioxide  with  the  oxygen  of 
the  air  blast.  But  the  dioxide  is  soon  reduced  by  the  hot 
carbon  to  the  monoxide,  which  interacts  with  the  ore,  thus:  — 

3Fe203      +         CO  2Fe3O4          +       CO2 

Ferric  Carbon  Ferrous-ferric  Carbon 

Oxide  Monoxide  Oxide  Dioxide 

Fe3O4     +     CO     =       3FeO     +     C02 

Ferrous 
Oxide 

FeO  +  CO  =  Fe  +  CO2 

At  this  stage  the  ore,  though  not  wholly  reduced,  becomes 
soft  and  porous,  and  as  the  mass  sinks  into  the  hottest  part 
of  the  furnace  the  reduction  is  completed,  thus  :  — 

FeO  +  C  =  Fe  4-  CO 

The  iron  now  combines  with  a  small  percentage  of  carbon, 
melts,  and  sinks  through  the  molten  slag.  This  iron  contains 
small  amounts  of  carbon,  silicon,  manganese,  phosphorus, 
and  sulphur. 

Varieties  of  Iron.  —  The  iron  we  use  and  speak  of  is  not 
pure  iron,  but  a  mixture  or  compound  of  iron  with  other 
elements,  chiefly  carbon.  It  is  customary  to  speak  of  three 
varieties  of  iron,  —  cast  iron,  wrought  iron,  and  steel. 
This  classification  is  based  chemically  upon  the  per  cent  of 
carbon  they  contain,  though  their  physical  properties  are 
also  modified  by  the  presence  of  silicon,  phosphorus,  sulphur, 


IRON  537 

and  manganese,  as  well  as  by  the  method  of  manufacture. 
The  different  varieties  are  closely  related,  and  pass  easily 
and  gradually  into  each  other.  Commercially,  there  are 
several  kinds  of  cast  iron  and  many  kinds  of  steel. 

Cast  Iron  is  the  most  impure  variety.  It  contains,  besides 
carbon,  the  impurities  already  mentioned.  The  carbon 
varies  from  3  to  5  per  cent,  the  silicon  and  manganese  are 
each  about  3  per  cent,  while  the  proportion  of  phosphorus  and 
sulphur  is  small.  If  molten  iron  is  cooled  suddenly,  the  prod- 
uct, which  is  very  brittle,  is  called  white  cast  iron ;  the  car- 
bon in  it  is  mostly  in  the  form  of  a  carbide  (cementite,  Fe3C). 
By  cooling  slowly,  much  of  the  carbon  remains  uncombined 
as  hard  crystals  known  as  graphite  carbon,  and  the  color  of 
the  iron  is  gray;  this  kind  is  gray  cast  iron.  It  is  softer 
than  the  white  variety,  and  melts  at  a  lower  temperature. 
Although  cast  iron  is  brittle,  it  will  withstand  great  pressure. 
Owing  to  its  crystalline  structure,  it  cannot  be  welded  or 
forged;  that  is,  hot  pieces  cannot  be  united,  nor  be  shaped 
by  hammering.  But  it  is  extensively  used  to  make  castings. 
This  is  the  kind  of  iron  used  in  an  ordinary  iron  foundry. 
The  iron,  which  melts  at  about  1200°  C.  (depending  upon  the 
impurities),  is  heated  in  a  furnace  similar  to  a  blast  furnace, 
and  when  molten  is  poured  into  sand  molds  of  the  desired 
shape.  Stoves,  pipes,  pillars,  railings,  parts  of  machines, 
and  many  other  useful  objects  are  made  of  cast  iron. 

Cast  iron  containing  5  to  20  per  cent  of  manganese  is  called 
spiegel  iron,  while  ferro-manganese  contains  from  20  to 
85  per  cent  of  manganese.  Both  are  used  in  making  steel. 

Wrought  Iron  is  the  purest  variety  of  commercial  iron. 
It  contains  not  more  than  0.5  per  cent  of  carbon  and  some- 
times only  0.06  per  cent,  the  average  being  0.15  per  cent. 
It  is  tough,  malleable,  and  fibrous,  and  can  be  bent.  Unlike 
cast  iron,  it  does  not  withstand  pressure,  but  it  will  sustain 


538  INORGANIC  CHEMISTRY 

great  weight.  An  iron  wire  will  sustain  the  weight  of 
nearly  a  mile  of  itself.  Wrought  iron  melts  at  such  a  high 
temperature  (1550°  to  2000°  C.)  that  it  is  not  used  for 
casting;  it  softens  at  a  relatively  low  temperature  (about 
1000°  C.),  can  be  forged  and  welded,  and  is  often  called 
malleable  iron.  It  may  be  seen  undergoing  these  opera- 
tions in  a  blacksmith's  shop.  It  can  also  be  rolled  into 
sheets  and  plates  and  drawn  into  fine  wire;  in  these  forms 
the  metal  is  very  tough.  Wrought  iron  is  made  into  wire, 
sheets,  rods,  nails,  spikes,  bolts,  chains,  anchors,  horse- 
shoes, tires,  and  agricultural  implements.  It  is  less  im- 
portant than  formerly,  since  it  is  being  largely  replaced  by 
steel. 

Wrought  iron  is  made  from  cast  iron  by  burning  out  most 
of  the  impurity.  Cast  iron  together  with  a  little  scrap  iron 
and  flux  is  melted  in  a  furnace,  much  like  a  reverberatory 
furnace,  lined  on  the  bottom  and  sides  with  iron  ore  (ferric 
oxide,  Fe2O3).  The  impurities  are  oxidized  partly  by  the 
oxygen  of  the  air,  but  mainly  by  the  oxygen  of  the  iron  oxide  ; 
the  silicon,  phosphorus,  sulphur,  and  manganese  pass  off 
in  the  slag ;  most  of  the  carbon  escapes  as  an  oxide,  though 
a  small  amount  remains  in  the  iron.  As  these  elements  are 
removed,  the  mass  becomes  pasty,  owing  to  the  higher 
melting  point  of  the  pure  iron.  It  is  now  stirred  vigorously, 
or  "  puddled."  At  the  proper  time  lumps  called  "  blooms" 
are  removed  and  hammered,  or  more  often  rolled  between 
ponderous  rollers.  This  operation  removes  the  slag,  and  if 
the  rolling  is  repeated,  the  quality  of  the  iron  is  improved; 
the  final  rolling  often  leaves  the  iron  in  the  shape  desired 
for  market. 

Steel  is  usually  intermediate  between  cast  iron  and  wrought 
iron  as  far  as  its  proportion  of  carbon  is  concerned.  Many 
grades  of  steel  are  manufactured,  and  their  physical  prop- 


IRON  539 

erties  do  not  depend  merely  upon  the  presence  of  a  small 
proportion  of  carbon  and  other  elements,  especially  phos- 
phorus, silicon,  and  certain  metals,  but  to  a  considerable 
extent  upon  the  method  of  manufacture  and  subsequent 
treatment.  Considered  from  the  standpoint  of  composition, 
it  may  be  said  that  in  general  the  higher  the  percentage  of 
carbon,  the  harder  the  steel.  In  soft  or  mild  steel  the  carbon 
is  seldom  more  than  .2  per  cent,  while  in  hard  steel  the 
amount  may  be  as  high  as  1.5  per  cent. 

Manufacture  of  Steel.  —  The  aim  in  the  manufacture  of 
steel  is  to  prepare  a  product  containing  little  or  no  sulphur, 
phosphorus,  and  silicon,  but  the  desired  proportion  of  carbon. 
This  is  accomplished  by  several  processes,  viz.  the  Bessemer, 
open-hearth,  crucible,  and  cementation. 

(1)  The  Bessemer  process,  which  is  quite  generally  used, 
was  devised  about  1860,  and  has  practically  revolutionized 
steel  making.  The  process  consists  in  burning  out  the 
impurities  in  cast  iron  by  forcing  air  through  the  molten 
metal,  and  then  adding  just  enough  iron  of  known  com- 
position to  give  the  desired  proportion  of  carbon.  The 
operation  is  carried  on  in  a  converter  (Fig.  81).  This  is 
a  huge,  pear-shaped  vessel,  supported  so  that  it  can  be 
tipped  into  different  positions;  it  is  also  provided  with 
holes  (C,  C,  C)  at  the  bottom,  through  which  a  powerful  blast 
of  air  can  be  blown.  It  is  made  of  thick  wrought-iron  plates, 
and  is  lined  with  an  infusible  mixture,  usually  rich  in  silica. 
The  converter  when  in  use  is  swung  into  a  horizontal  position 
(A),  and  five  to  twenty  tons  of  molten  pig  iron  are  poured 
in,  often  directly  from  the  blast  furnace.  The  air  blast  is 
turned  on  and  the  converter  is  swung  back  to  a  vertical 
position  (B).  As  the  air  is  forced  in  fine  jets  through  the 
molten  metal,  the  temperature  rises,  and  the  carbon,  silicon, 
and  manganese  are  oxidized.  The  carbon  forms  carbon  mon.- 


540 


INORGANIC  CHEMISTRY 


oxide,  which  burns  at  the  mouth  of  the  converter,  while  the 
other  oxides  pass  into  the  slag.  .  This  oxidation  generates 
enough  heat  to  keep  the  metal  melted,  and  no  fuel  need  be 
used.  As  soon  as  the  impurities  have  been  burned  out, 
sufficient  spiegel  iron  or  ferromanganese  is  added  to  furnish 
the  proper  amount  of  carbon  and  manganese.  By  adding 


Air- 


FIG.  81.  —  Converter. 


spiegel  iron  of  known  composition,  Bessemer  steel  of  any 
desired  grade  is  produced.  After  the  completion  of  the 
operation,  which  takes  about  twenty  minutes,  the  metal  is 
poured  from  the  converter  into  molds  to  cool.  These  cold 
blocks  of  steel  are  called  ingots. 

In  the  Bessemer  process,  as  described  in  the  preceding 
paragraph,  sulphur  and  phosphorus  are  not  removed  from 
the  steel.  Both  are  objectionable  impurities;  sulphur  makes 
steel  brittle  when  hot,  and  phosphorus  makes  it  brittle  when 
cold.  The  Thomas-Gilchrist  process  (or  basic  process)  is 
a  modification  of  the  Bessemer  process  by  which  the  sulphur 
and  phosphorus  can  be  removed  in  the  slag.  The  con- 
verter in  this  process  is  lined  with  burned  dolomite  (i.e. 
practically  a  mixture  of  lime  and  magnesia),  called  a  basic 
lining ;  lime  is  also  added  to  the  charge  of  pig  iron,  and  the 


IRON 


541 


air  blast  is  continued  a  little  longer  than  in  the  Bessemer 
process;  otherwise  the  operations  are  the  same.  The  car- 
bon passes  off  as  usual,  and  the  oxides  of  phosphorus,  silicon, 
and  sulphur  combine  with  the  basic  constituents  of  the  lining 
to  form  a  slag.  This  lining,  after  use,  is  known  as  Thomas 
slag;  it  is  utilized  as  a  fertilizer  on  account  of  its  phosphorus 
content. 

(2)  In  the   Siemens-Martin   or   open-hearth   process   cast 
iron,  or  a  mixture  of  cast  iron,  scrap  iron,  and  iron  ore,  in 


FIG.  82.  —  Open-hearth  furnace. 

proper  proportions,  is  melted  by  hot  gases  in  a  special 
kind  of  furnace  called  an  open-hearth  furnace  (Fig.  82). 
The  receptacle,  or  hearth  ( H) ,  in  which  the  charge  is  melted, 
is  lined  with  sand  in  the  acid  process  or  with  lime  and  magne- 
sia in  the  basic  process,  thereby  permitting  the  removal  of 
all  the  impurities.  At  the  base  of  the  furnace,  are  duplicate 
sets  of  checkerwork  (A,  B  and  .C,  D).  As  the  hot  gases 
pass  through  A,  B  to  the  chimney,  they  heat  the  checker- 
work.  The  fuel  gas  is  then  passed  through  B  and  air 
through  A.  The  mixture  of  air  and  gas  burns  and  pro- 


542  INORGANIC  CHEMISTRY 

duces  a  much  higher  temperature  on  the  hearth  than  if 
the  gaseous  mixture  were  cool.  The  oxidizing  flame  passes 
over  the  charge  on  the  hearth  (H),  oxidizing  some  of  the 
impurities  and  keeping  the  mass  at  such  a  temperature 
that  other  impurities  form  a  slag  with  the  lining.  Mean- 
while the  hot  products  of  combustion  and  the  unused  gases 
are  passed  through  the  checkerwork  C,  D  and  heat  it. 
The  fuel  gas  and  air  are  then  made  to  pass  through  this 
checkerwork  to  the  hearth  and  out  over  the  other  checker- 
work  (A,  B)  to  the  chimney.  Thus  the  process  is  alternated, 
one  checkerwork  being  cooled  as  the  other  is  heated,  and 
vice  versa.  It  is  only  by  this  regenerative  process,  as  it  is 
called,  that  enough  heat  is  obtained  to  keep  the  charge 
melted  as  it  becomes  purer  and  purer.  The  charge  is  heated 
from  eight  to  ten  hours ;  when  a  test  shows  that  the  metal 
has  the  desired  composition,  certain  metals  or  ferro-alloys 
are  added,  and  the  steel  is  quickly  poured  into  molds  and 
allowed  to  cool  into  ingots.  Subsequently  the  ingots  of  this 
kind  of  steel  (as  well  as  of  other  kinds)  are  reheated,  rolled, 
and  cut  into  lengths  of  the  proper  size  and  shape,  being 
then  known  as  billets.  The  open-hearth  process  requires 
a  special  furnace  and  gas  plant,  and  is  more  expensive  than 
the  Bessemer  process,  since  it  takes  much  longer.  But  it 
is  easily  controlled,  and  yields  a  tough,  elastic  steel,  which 
is  excellent  for  bridges,  large  machines,  large  guns,  and  gun 
carriages.  The  production  of  the  open-hearth  steel  has 
increased  rapidly  in  the  last  few  years. 

(3)  In  the  crucible  process  wrought  iron  is  melted  with 
charcoal  in  graphite  or  clay  crucibles.     During  the  melting, 
which  lasts  three  or  four  hours,  the  iron  is  slowly  changed 
into  steel  by  absorbing  the  proper  proportion  of  carbon. 

(4)  In  the  cementation  process  wrought  iron  and  carbon 
are  heated  in  fire-brick  boxes  for  several  days.     The  trans- 
formation is  the  same  as  in  the  crucible  process. 


IRON  543 

The  four  processes  of  making  steel,  just  described,  provide 
many  grades  which  permit  its  use  for  countless  purposes. 
In  recent  years  special  steels  have  been  made  by  adding 
to  steel  small  quantities  of  other  metals,  such  as  nickel, 
chromium,  molybdenum,  tungsten,  vanadium,  and  manganese. 
Such  steels  differ  somewhat  in  special  properties,  but  all  are 
characterized  by  extreme  hardness,  toughness,  and  strength. 
The  metals  are  sometimes  added  directly  to  the  molten  steel, 
but  more  often  in  the  state  of  an  alloy  of  iron.  These 
alloys  contain  varying  but  known  percentages  of  the  con- 
stituents, and  are  called  ferrochrome,  ferrosilicon,  ferrotung- 
sten,  etc. 

The  important  properties  of  steel  are  numerous.  It  is 
fusible  and  malleable,  and  can  be  forged,  welded,  and  cast. 
It  is  harder,  stronger,  and  more  durable  than  pure  iron,  and 
is  therefore  more  serviceable.  But  its  most  valuable  prop- 
erty is  the  varying  hardness  which  it  may  be  made  to 
acquire.  If  steel  is  heated  very  hot  and  then  suddenly  cooled 
by  immersion  in  cold  water  or  oil,  it  becomes  brittle  and  very 
hard.  But  if  heated  and  cooled  slowly,  it  becomes  soft, 
tough,  and  elastic.  All  grades  of  hardness  may  be  obtained 
between  these  extremes.  And  if  the  hardened  steel  is  re- 
heated to  a  definite  temperature,  determined  approximately 
by  the  color  the  oxidized  metal  assumes,  and  then  properly 
cooled,  a  definite  degree  of  hardness  and  elasticity  is  ob- 
tained. This  last  operation  is  called  tempering. 

Uses  of  Steel.  —  Steel  is  now  used  instead  of  iron  for  many 
purposes.  High  buildings,  bridges,  rails,  cars,  locomotives, 
battleships,  electrical  machinery,  boilers,  agricultural  im- 
plements, wire  nails,  rods,  hoops,  tin  plates,  and  castings 
of  all  kinds  consume  vast  amounts  of  Bessemer  and  open- 
hearth  steel.  Crucible,  cementation,  and  special  steels  are 
used  in  making  springs,  tools,  cutlery,  pens,  and  needles. 


544  INORGANIC  CHEMISTRY 

Properties  of  Pure  Iron.  —  Chemically  pure  iron  can  be 
obtained  as  a  black  powder  by  reducing  the  oxide  with 
hydrogen.  Recently  very  pure  massive  iron  has  been  pre- 
pared electrolytically  from  a  solution  of  ferrous  and  am- 
monium sulphates.  The  purest  commercial  form  is  the 
wrought  iron  used  for  piano  wire.  Pure  iron  is  a  silvery 
white,  lustrous  metal.  It  is  softer  than  ordinary  iron,  but 
melts  at  a  higher  temperature  (about  1520°  C.).  The 
specific  gravity  is  about  7.8.  It  is  attracted  by  a  magnet, 
but  soon  loses  its  own  magnetism.  Dry  air  has  no  effect 
upon  iron,  but  moist  air  rusts  it.  The  rusting  of  iron  is  a 
complex  process.  Interpreted  by  the  electrolytic  dissociation 
theory  the  iron  first  interacts  with  water  and  goes  into  solu- 
tion as  ferrous  ions  (Fe"1"1"),  while  the  hydrogen  ions  (H+)  be- 
come hydrogen  molecules  and  escape ;  the  ferrous  ions  com- 
bine with  the  hydroxyl  ions  (OH~)  left  in  the  water  and  form 
ferrous  hydroxide  (Fe(OH)2),  which  subsequently  becomes 
iron  rust.  Rusting  proceeds  rapidly,  because  the  film  of  rust 
is  not  compact  enough  to  protect  the  metal.  Iron  readily 
interacts  with  dilute  acids,  and  as  a  rule  hydrogen  and  ferrous 
compounds  are  the  products. 

With  nitric  acid  various  products  result,  according  to  the  con- 
ditions, —  ferrous  nitrate  and  ammonium  nitrate  if  the  acid  is 
cold,  but  ferric  nitrate  and  oxides  of  nitrogen  if  the  acid  is  warm. 
If  a  clean  iron  wire  is  dipped  into  fuming  nitric  acid  and  then  into 
ordiaary  nitric  acid,  no  action  is  apparent.  The  iron  is  said  to  be 
passive.  This  peculiar  fact  has  not  been  adequately  explained. 
Steam  and  hot  iron  interact  thus :  — 

3Fe    +    4H2O    =        Fe3O4       +        4  H2 

Iron  Water  Iron  Oxide  Hydrogen 

The  film  of  the  oxide  (FesCU),  unlike  iron  rust,  adheres  firmly  and 
protects  the  iron  from  further  oxidation. 

Compounds  of  Iron.  —  Iron  forms  two  important  series 
of  compounds,  —  ferrous  and  ferric.  They  are  analogous 


IRON  545 

to  cuprous  and  cupric,  mercurous  and  mercuric  compounds. 
The  valence  of  iron  is  two  in  ferrous  compounds  and  three 
in  ferric.  Ferrous  compounds  in  an  acid  solution  pass  into 
the  corresponding  ferric  compound  by  the  action  of  oxidiz- 
ing agents,  e.g.  oxygen,  nitric  acid,  potassium  chlorate, 
potassium  permanganate,  and  chlorine.  Conversely,  ferric 
compounds  are  reduced  to  the  ferrous  by  reducing  agents, 
e.g.  hydrogen,  hydrogen  sulphide,  sulphur  dioxide,  and 
stannous  chloride.  The  passage  from  one  series  to  the  other 
occurs  easily,  especially  from  ferrous  to  ferric.  The  oxida- 
tion and  reduction  of  iron  compounds  illustrates  typically 
the  broad  use  of  these  terms.  Oxygen  is  not  necessarily 
involved.  Thus,  ferrous  and  ferric  chlorides  pass  readily 
into  each  other  by  the  addition  or  removal  of  chlorine.  The 
two  processes  are  general  and  mutual,  and  may  be  sum- 
marized as  follows:  (a)  Oxidation  is  the  chemical  addition 
of  oxygen  or  any  other  negative  element,  such  as  chlorine; 
reduction  is  the  removal  of  oxygen  or  any  other  negative 
element,  (b)  From  the  standpoint  of  the  ionic  theory, 
oxidation  is  the  addition  of  positive  electricity,  and  reduc- 
tion the  withdrawal.  Thus,  when  ferrous  chloride  solution 
becomes  ferric  chloride  upon  the  addition  of  chlorine, 
the  ferrous  ions  (Fe++)  become  ferric  ions  (Fe+++),  while 
the  electrically  neutral  chlorine  molecules,  having  become 
ions,  lose  positive  electricity,  i.e.  they  undergo  reduction, 
(c)  Occasionally,  oxidation  is  spoken  of  as  occurring  when 
the  valence  is  raised,  and  reduction  when  the  valence  is 
lowered.  For  instance,  in  the  change  from  ferrous  to  ferric 
compounds  the  valence  of  iron  is  raised  from  two  to  three, 
while  by  the  reduction  of  ferric  to  ferrous  the  valence  is  low- 
ered from  three  to  two.  (Compare  interaction  of  mercuric 
chloride  and  stannous  chloride.) 

Oxides  and  Hydroxides  of  Iron.  —  Iron  forms  three  oxides. 


546  INORGANIC  CHEMISTRY 

Ferrous  oxide  (FeO)  is  an  unstable  black  powder.  Ferric 
oxide  (Fe2O3)  occurs  native  in  many  varieties  as  hematite 
—  the  most  abundant  ore  of  iron.  It  can  be  prepared  by 
heating  ferrous  sulphate  or  ferric  hydroxide.  Large  quan- 
tities are  manufactured  from  the  ferrous  sulphate  obtained 
as  a  by-product  in  the  cleansing  of  the  iron  used  in  making 
galvanized  and  tinned  ware.  It  is  sold  under  the  names 
rouge,  crocus,  and  Venetian  red,  and  is  used  to  polish  glass 
and  jewelry  and  to  make  red  paint.  Ferrous-ferric,  or 
ferroso-ferric,  oxide  (magnetic  oxide  of  iron,  Fe3O4),  occurs 
native  as  magnetite;  if  magnetic,  it  is  called  loadstone. 
It  is  produced  as  a  film  or  scale  by  heating  iron  in  the  air. 
The  firm  coating  of  this  oxide  formed  by  exposing  iron  to 
steam  protects  the  metal  from  further  oxidation ;  iron  thus 
coated  is  called  Russia  iron.  Some  authorities  call  this  oxide 
iron  ferrite  (Fe(FeO2)2).  Ferrous  hydroxide  (Fe(OH)2)  is  a 
white  solid  formed  by  the  interaction  of  a  ferrous  salt  and  an 
alkali,  such  as  sodium  hydroxide.  Exposed  to  the  air,  it  soon 
turns  green,  and  finally  brown,  owing  to  the  formation  of 
ferric  hydroxide.  Ferric  hydroxide  (Fe(OH)3)  is  a  reddish 
brown  solid,  formed  by  the  interaction  of  an  hydroxide  (e.g. 
sodium  hydroxide)  and  a  ferric  salt.  It  readily  forms  a 
colloidal  solution  (compare  Arsenic  Trisulphide,  page  404). 

Ferrous  Sulphate,  FeS04,  is  a  green  salt  obtained  by  the 
interaction  of  iron  (or  ferrous  sulphide)  and  dilute  sulphuric 
acid,  and  is  a  by-product  in  several  large  industries  (e.g. 
see  Ferric  Oxide).  It  is  also  prepared  on  a  large  scale  by 
oxidizing  iron  pyrites  (FeS2) .  This  is  accomplished  simply  by 
roasting,  or  more  often  by  exposing  heaps  of  pyrites  to  moist 
air;  the  mass  is  extracted  with  water  containing  scrap  iron 
and  a  small  proportion  of  sulphuric  acid,  and  large  light 
green  crystals  are  obtained  from  the  solution.  The  crys- 
tallized salt  (FeS04 . 7  H2O)  is  also  called  green  vitriol  or 


IRON  547 

copperas.  Exposed  to  the  air,  ferrous  sulphate  effloresces 
and  oxidizes.  Large  quantities  are  used  as  a  mordant  in 
dyeing  silk  and  wool,  as  a  disinfectant  and  a  wood  preserva- 
tive, and  in  manufacturing  ink,  bluing,  pigments,  leather, 
varnish,  and  mottled  soaps.  Much  black  writing  ink  is 
made  essentially  by  mixing  ferrous  sulphate,  nutgalls,  gum, 
and  water.  A  mixture  of  lime  and  ferrous  sulphate  is  used 
as  a  coagulant  in  purifying  water  and  sewage. 

Ferric  Sulphate,  Fe2(SO4)3,  is  formed  by  oxidizing  an 
acid  solution  of  ferrous  sulphate  with  nitric  acid.  When 
ferric  sulphate  solution  is  mixed  with  the  proper  quan- 
tity of  potassium  (or  ammonium)  sulphate,  iron  alum 
(K2SO4 .  Fe2(SO4)3 .  24  H2O  or  K2Fe2(S04)4 .  24  H20)  is  formed. 
It  is  a  pale  violet,  crystalline  solid,  which  has  properties 
like  ordinary  alum.  Iron  alum  is  used  chiefly  as  a 
mordant. 

Iron  Sulphides.  —  There  are  two  iron  sulphides.  Com- 
mercial ferrous  sulphide  (FeS)  is  a  black,  brittle,  metallic- 
looking  solid,  but  the  pure  compound  is  yellow  and  crys- 
talline. It  is  also  obtained  as  a  fine  black  precipitate  by 
the  interaction  of  a  dissolved  ferric  or  ferrous  salt  and  am- 
monium sulphide,  though  not  by  hydrogen  sulphide.  It  is 
made  on  a  large  scale  by  fusing  a  mixture  of  iron  and  sulphur. 
Its  chief  use  is  in  preparing  hydrogen  sulphide.  Ferric 
sulphide  (iron  disulphide,  iron  pyrites,  pyrite,  FeS2)  is  one 
of  the  commonest  minerals.  It  is  a  lustrous,  metallic, 
brass-yellow  solid.  Crystals  of  pyrites,  found  in  many  rocks, 
are  often  mistaken  for  gold  —  hence  the  popular  name, 
"  fool's  gold."  It  is  valueless  as  an  iron  ore,  but  large  quan- 
tities are  used  as  a  source  of  sulphur  in  making  sulphuric 
acid.  Over  one  and  a  half  million  tons  are  annually  con- 
sumed in  the  sulphuric  acid  industry. 


548  INORGANIC  CHEMISTRY 

Iron  Chlorides.  —  When  iron  interacts  with  hydrochloric 
acid,  ferrous  chloride  (FeCl2)  is  formed  in  solution.  Ferric 
chloride  (FeCl3)  is  readily  prepared  by  passing  chlorine  gas 
into  a  solution  of  ferrous  chloride.  It  is  a  dark,  lustrous, 
crystalline  solid ;  but  owing  to  its  extreme  deliquescence, 
it  is  usually  sold  as  a  solution,  which  is  a  brown  liquid.  Nas- 
cent hydrogen  or  another  reducing  agent  changes  ferric  chlo- 
ride into  ferrous  chloride.  It  hydrolyzes  readily. 

Ferrous  Carbonate,  FeCOs,  occurs  native  as  the  iron  ore 
siderite,  clay  ironstone,  or  spathic  iron  ore.  The  typical 
variety  is  light  yellow  or  brown,  lustrous,  crystalline,  and 
not  very  hard;  but  many  kinds  are  impure,  and  the  prop- 
erties vary.  It  is  slightly  soluble  in  water  containing 
carbon  dioxide,  and  is  therefore  found  in  some  mineral 
springs.  (See  Chalybeate  Waters.)  Like  all  carbonates, 
it  yields  carbon  dioxide  with  warm  hydrochloric  acid. 

Iron  Cyanides.  —  Iron  and  cyanogen  (CN)2,  with  or  with- 
out potassium,  form  several  compounds.  The  most  im- 
portant is  potassium  ferrocyanide  (K4Fe(CN)6).  It  is  a 
lemon-yellow,  crystalline  solid,  containing  three  molecules 
of  water  of  crystallization.  Unlike  most  cyanogen  com- 
pounds, it  is  not  poisonous.  Its  commercial  name  is  yellow 
prussiate  of  potash.  It  is  manufactured  by  fusing  together 
iron  filings,  potassium  carbonate,  and  nitrogenous  animal 
matter  (such  as  horn,  hair,  blood,  feathers,  and  leather). 
The  mass  is  extracted  with  water,  and  the  salt  is  separated 
by  crystallization.  In  Germany  this  salt  is  manufactured 
from  the  iron  oxide  which  has  been  used  to  purify  illuminat- 
ing gas.  Large  quantities  are  used  in  dyeing  and  calico 
printing,  and  in  making  bluing  and  potassium  cyanogen 
compounds.  Potassium  ferricyanide  (K3Fe(CN)6)  is  a  dark 
red,  crystalline  solid,  containing  no  water  of  crystallization. 
It  is  often  called  red  prussiate  of  potash.  It  is  manufactured 


IRON  549 

by  oxidizing  potassium  ferrocyanide  with  potassium  per- 
manganate or  chlorine,  thus:  — 

2  K4Fe(CN)6     +     C12     =     2  K3Fe(CN)6     +     2  KC1 

Potassium  Chlorine  Potassium  Potassium 

Ferrocy?nide  Ferricyanide  Chloride 

It  is  very  soluble  in  water,  forming  a  yellow,  unstable  solu- 
tion. In  alkaline  solution  it  is  a  vigorous  oxidizing  agent, 
and  finds  extensive  use  in  dyeing.  It  is  also  used  as  one  of 
the  ingredients  of  the  sensitive  coating  of  "  blueprint"  paper. 

The  valence  of  the  radical  Fe(CN)6  is  three  in  ferri- 
cyanides  and  four  in  ferrocyanides. 

Ferrous  salts  and  potassium  ferricyanide  interact  in  solu- 
tion and  produce  ferrous  ferricyanide  (Fe3(Fe(CN)6)2). 
This  is  a  blue  solid  and  is  often  called  Turnbull's  blue.  But 
ferrous  salts  produce  with  potassium  ferrocyanide  a  white 
precipitate  (ferrous  ferrocyanide)  which  quickly  oxidizes 
to  a  complex  blue  compound.  Ferric  salts  interact  with 
potassium  ferrocyanide  and  produce  ferric  ferrocyanide 
(Fe4(Fe(CN)6)3).  This  is  a  dark  blue  solid,  and  is  called 
Prussian  blue  or  Berlin  blue.  Ferric  salts  produce  no  pre- 
cipitate with  potassium  ferricyanide.  Prussian  blue  is  ex- 
tensively used  in  dyeing  and  calico  printing,  and  in  making 
bluing.  The  above  reactions,  which  allow  ferrous  and 
ferric  salts  to  be  distinguished,  may  be  summarized  as 
follows :  — 


CYANIDE 

FERROUS    SALT 

FERRIC  SALT 

Ferrocyanide 
Ferricyanide 

Whitish  precipitate 
Turnbull's  blue 

Prussian  blue 
No  precipitate 

Besides  the  above  tests,  potassium  sulphocyanate  (KCNS) 
produces  a  blood-red  solution  of  ferric  sulphocyanate 
(Fe(CNS)3)  with  ferric  salts,  but  leaves  ferrous  salts  un- 
changed. The  tests  for  iron  are  thus  numerous  and  specific. 


550  INORGANIC  CHEMISTRY 

Miscellaneous.  —  Iron  compounds  yield  several  kinds  of 
ions,  e.g.  the  ferrous  ion  (Fe++)  and  the  ferric  ion  (Fe4"^"1")  ; 
each  is  colorless,  though  the  former  appears  delicate  green 
and  the  latter  very  pale  yellow,  owing  to  traces  of  other  sub- 
stances. Complex  ions  are  Fe(CN)6  and  Fe(CN)6 . 

The  atomic  weight  of  iron  is  55.84. 

NICKEL 

Nickel,  Ni,  occurs  combined  with  arsenic  or  with  arsenic 
and  sulphur  as  niccolite  (NiAs)  and  nickel  glance  (NiAsS). 
Small  amounts  of  metallic  nickel  are  found  in  meteorites. 
The  ores  which  furnish  most  of  the  commercial  nickel  are 
the  nickel-bearing  iron  sulphides  of  the  Sudbury  district, 
Canada,  and  the  silicate  of  nickel  and  magnesium  (gar- 
nierite)  found  in  New  Caledonia. 

Preparation  and  Properties.  —  Nickel  is  obtained  from  its 
ores  by  a  complicated  smelting  or  electrolytic  process.  It 
is  a  white,  hard  metal.  It  is  ductile  and  tenacious.  It 
takes  a  brilliant  polish  and  does  not  tarnish  in  dry  air, 
though  in  moist  air  it  tarnishes  very  slowly.  Like  iron,  it 
is  attracted  by  a  magnet. 

Uses  of  Nickel.  —  For  many  years  it  has  been  used  as  one 
ingredient  of  the  small  coins  of  several  countries.  The  per 
cent  of  nickel  varies  from  12  in  the  United  States  cent  to 
25  in  the  five-cent  piece.  German  silver  contains  from  15 
to  25  per  cent  of  nickel,  the  rest  being  copper  and  zinc. 
Large  quantities  of  nickel  are  used  to  coat  or  plate  other 
metals,  especially  iron  and  brass.  The  nickel  plating  is 
done  by  electrolysis,  as  in  the  case  of  silver  and  gold  plating, 
though  the  electrolytic  solution  used  is  a  sulphate  of  nickel 
and  ammonium  ((NH4)2SO4.  NiSO4)  — not  a  cyanide,  as  in 
the  other  cases.  The  deposit  of  nickel  is  hard,  brilliant, 


NICKEL  551 

and  durable.  Nickel  becomes  malleable  if  a  little  magnesium 
or  aluminium  is  added  to  the  molten  metal,  and  sheets  of 
iron  covered  with  such  nickel  are  made  into  household 
utensils.  Nickeloid  is  a  nickel-plated  sheet  zinc.  Its 
attractive  appearance  and  non-corrosive  property  adapt  it 
for  the  manufacture  of  reflectors,  refrigerator  linings,  bath 
tubs,  show  cases,  and  signs.  Nickel  is  used  in  the  manu- 
facture of  nickel  steel.  This  contains  varying  proportions 
of  nickel ;  large  quantities  are  used  for  burglar-proof  safes, 
and  the  armor  plates  and  turrets  of  battleships.  An  alloy 
of  nickel  and  copper  known  as  monel  metal  is  used  where  a 
non-corroding  metal  is  desirable. 

Compounds  of  Nickel.  —  Nickel  forms  two  series  of  com- 
pounds —  the  nickelous  and  the  nickelic.  The  valence  of 
nickel  is  two  in  the  nickelous  compounds  and  three  in  the 
nickelic.  The  nickelous  compounds  are  the  more  common. 
Many  nickel  salts  are  green.  The  apple-green  nickelous 
hydroxide  (Ni(OH)2)  is  formed  by  the  interaction  of  an 
alkali  and  a  dissolved  nickel  salt.  Nickelous  sulphide  (NiS) 
is  obtained  as  a  black  precipitate  by  the  interaction  of 
ammonium  sulphide  and  a  dissolved  nickel  salt.  Nickelous 
sulphate  (NiSO4)  and  nickelous  chloride  (NiCl2)  are  the  usual 
commercial  salts.  Nickel  carbonyl  (Ni(CO)4)  is  prepared 
by  passing  carbon  monoxide  over  finely  divided  nickel. 
It  is  a  volatile,  colorless  liquid  which  boils  at  43°  C.  The 
vapor  is  poisonous  and  decomposes  at  150°-180°  C.  into 
metallic  nickel  and  carbon  monoxide.  This  compound  has 
a  technical  application  in  the  Monde  process  of  preparing 
nickel.  Ammonium  nickelous  sulphate  ((HN4)2SO4 .  NiS04) 
was  mentioned  in  the  preceding  paragraph. 

Tests  for  Nickel.  —  The  formation  of  the  green  nickelous 
hydroxide  is  a  distinctive  test.  Nickel  compounds  color 


i\<>i«;.\Mr  <  MI.MISTKY 
a  borax  bead  brown  in  the  oxidizing  flame  and  gray  in  the 

reducing    flame. 

MilcelUneoui,  —  NiokelouH  Halts  yield  a  p-een  ion  (Ni1  '). 
wi'i/d.  :    •  I.  el   i.:  .f.s.(i.s. 


Cobalt,    ('<>,    flOnorally    OCOlirs    cumhined     willi     arsenic    or 
:n    emr      ;in.  I     .sulphur     ns      smallile      «'..\,;,)      and     cohallile 
(CoAsS),    which   JUT  u.:iiall\    :i     ocialed    \\illi    NIC  correspond 
ing    llirl.cl    r.Mii|Mi||||(lM. 

Preparation  and   Properties.  —  Cobalt  in   ohininnl   as  a 

powder  l.y  icdiiciii,^  il  nxidr  \\illi  hydrogen  or  :i  ;i  c<  .licrrnl, 
mass  I  iv  MIC  .MluiiiiiioMii'i  inic  mcMiiHl.  h  is  :i  \\iuh-  inctiil 
\\iMi  ;i  l:iinl.  reddish  Uli^o.  Like  nickel,  il  |,ul\(»H  :i  l>nlli:ml. 
polioh.  It  IH  Irss  ni.'i^tirlic  than  iron.  \l.  i.illic  colcdl.  has 
foW  UHOH, 

Compound!  of  Cobalt,       ('..!..  -ill,  liko  nickel,   forms  Two 

sci'irs  i.l'  coinpoiinds  Mir  rohalloiis  :in.|  MIC  c<  >l  i:ill  ic.  The 
v.'ilence  n|'  rnli:dl  is  I  \\  ......  'ohalioilM  r«  n  1  1  p<  n  1  1  id;  :  :ind  Miree 

III     Col  Xllt'lC.          I   lie     Col  i.'ll  I  OIIS    ci  ill  |  pi  ill  ||(  I        :iie     Mie     MHHC     e«iln 

moll,    Ihniiidi    ninny    complex    eolcill    n  unpi  >n  nd:;    ;iie    Kimwn. 

Cobaltoui  sulphide  (CoH)    is   nlit.-iini  .|    :i  i   .•>    M.-iek    pn-,-i|>i 

l;i!e    |»y     I  he     llilei  :ic!  IMII     of     ;|  1  1  1  1  1  1  Ml  1  1  1  II  1  1    ::i||phnle    :ilii|     H    dis- 

solved  cobalt  salt,  Cobaltous  chloride  (CoCl,)  and  cobaltoui 
nitrate  (Oo(NOa)i)  are  red  orystalline  salts;  they  orystnih  ••• 

\\Mli  six  iinileciilr::  uf  \\;iler  of  cry  ::l  :i  1  1  1  /;i  I  imi  .  WluMl  parl. 
or  all  ..f  Mie  wnler  ..f  crysl  alli/al  ion  is  driven  ..IT,  Miese  salts 
heroine  violet  or  Mile.  A  complex  cohalt  roni|»oiind  known 
OS  Smalt,  or  small  Mile,  i;  n  ,-,|  |«>  decorate  poie.-l.-.in.  1  1. 
h:i  a  \  anahle  coinpo  ilion,  lull  is  e  isenlially  a  cohnll  ilicale. 
Cobalt  Compounds  Color  glass  hlue.  Such  j-.lass  Iran  '11111  ; 
red,  Mne,  and  \  io|el.  lij^hl,  Mil  not  yellow,  orange,  and  p'een, 


COBALT  :>:,:; 

,sn   il    is  sumelimes   Used    In  deleel    I  lie   vinlH.   |  ><  >l  :i:  ;  I  HIM    ll.'ime 
\\  hen  in;i,;ked   l»v   I  In-  \ellu\\   sn( 


IVsIs  loi  Col>:dl.  ('nlcil!  eMinpnmids  cnlnr  n,  bnrux  IxMld 
Mm-  in  Imili  limn  (  'nkillmi;;  rnmpouridH  when  mixed 

\\iili    |H,I.I      mm    nil  rile     MM!    :ieelie   ;ieid    I'nrin    ;i    yellnw  while 

l>i-eei|iii;ih'  of  potassium  cobtltinitrite  (K»(  'm  N(  >•>,,>  Iherehy 
disl  i  n^iiish  ine  ••••!. .-ill  limn  nickel. 

|'U«IMI.|.;M:;    AMI.    K  \  KIM  'IMIOH 

1.  \\  h.il    JM   Mi-       ini|.le:;|,  formula  <»!'  n  compound   1)  gill,   of   which 
•   n   l<|.  <l     IS  fin.    «.l       iilphui    .unl     I   '.',  HIM.   of   iron    ' 

2.  (")   AII  iron  oxi<l<    '"ni.-i.iiiH  '27.0  IMT  cent  of  oxygon  and  IIUH 
Mm  iiiolcciil.u  \\<  i!-lil,  of  2!t2.     Wlial/  !H  il-H  HiiuploNl.  formula?     (/"    M 
I.OI.'i  K"'    of  iron  form   I.I  -Hi  jfm.  of  IMI  iron  uxido,  whnl   in  Mm  nlm- 

|»lc::l    r<>riiiill;i    ul'    I  lie   i»\|i|«     ' 

.'{.     (  ':ilciil:i!r    MM-    approximal.(i    M,toUli<'    \v« -ii-lil.    of    iron    fniiii    Mm 

|n    ,     |||,        I,,    ,|   I      ,,|       !    Ill        MM     I.I  I 

I        \\  li;il    \\rii-lil    "I    ii  -.  •,  ..i  n  i     t.  -i  i  ii  IK  .1   I"  ii  111  !<•  will)  21  ^111.  of  1 1"  1 1 

to  KIV«  1 1"  in.r'ii.  i  i<«  ..  i«l«  of  HOD  ?  Wlial,  volume  will  Miin  oxygon 
occupy  :il  In  (  '.  mid  /.>()  nun.  ' 

.'».      \\h;i!     \\<i"li!     of    cr\-:;|;illl/.co!     :i  II I  inoi  1 1 1 1  n  i     (foil     nJlllll      will     I  in 

formed  l>\  l  IK  mil  r.i<  1 1011  of  soluliotiH  condiiniiiK  12  KIM.  of  ammo- 
nium :;ill|.li;ilc  ;iud  .'{()  |>m  of  frrrir  :;nlplui  Ir  •' 

(i.  I  low  Illllcli  co|..c  (  coll  I  ;i  I  ii  in;-  'Mi  .1  |»cr  ecu  I  ol'  «  .illioin  I  1 1  ceded 
lo  reduce  I7f>  loii  of  li«  111. i  Ml-  I'M,  ,  per  ceiil  |IIIM  <  ' 

7.  <  ';ilcul:i  Ic  lln  \\«i"hl  ol  coli:ill  or  iiid.<l  in  (<i\  '»  "in  of 
poUiHsium  col»!ill  mil  rile,  l\  ('o(NO  In;  (/>)  •'  r.ni  of  colw.lfoiiH 
nil.ni,l,n;  (r)  KM)  millii'r.inr  -.1  nickel  carhonyl,  NiK'O),,  i./i  •.'::.. 
c.(  n  i  CM  . i  m:  of  :i  uiiiiouium  nickeloiiH  Nil  I  phal  <\  (Nlli).S(>,  .  NiH()<. 

H.    Wril.n  Mm  formuliiH  ol    MM    following  oompouud  i  ;uid  iudicnln 

MIC     '..I,  n,-,      of    e.'icll     clciiielll   :      I'olnHsillHl     I'd  i  < »      .1 1 1  n  I-  .     polllHHilllM 

fen  icyniiide,  ferrin  ferrocynuidc,  I'nrrou  IVrric  iiuidn,  poljiHHium 
eohall  iuil file. 

0.  I  low  i  n  mi  y  ft-,  of  ammonia  HO!  it  Mo  1 1  IUIVJIIK  a  Hpecillc  ft  .1  1 1  <>l 
.IMlOf)  and  containing  O.HM  per  <'en!,  of  Nil,,  li.v  w«  c-hi  an  im-dnd  to 
|.r.  -  -i|.il:ile  MM-  iron  ;,  l-'.-i «  )|  h  ,  from  I  r.m.of  (  N  1 1  <),S<  )4  .  KeHO4  , 
(.11,0? 


CHAPTER  XXXIV 
Platinum  and  Associated  Metals 

Occurrence  of  Platinum.  —  Platinum  occurs  as  the  essen- 
tial ingredient  of  platinum  ore  or  so-called  native  platinum. 
The  ore  contains  from  60  to  86  per  cent  of  platinum.  The 
other  metals  present  are  ruthenium,  osmium,  iridium, 
rhodium,  and  palladium.  Iron,  gold,  and  copper  are  also 
usually  present.  Only  one  native  compound  is  known, 
viz.  platinum  arsenide  (sperrylite,  PtAs2).  These  metals, 
like  gold,  are  noble  metals,  i.e.  they  do  not  unite  with  oxygen 
at  any  temperature. 

The  word  platinum  is  derived  from  platina,  a  form  of  the  Spanish 
word  plata,  meaning  silver,  because  native  platinum  was  regarded 
as  an  impure  ore  of  silver  by  the  Spaniards,  who  first  discovered  it 
in  South  America  about  1735.  Platinum  is  now  sometimes  called 
by  its  old  name  platina. 

Preparation  of  Platinum.  —  The  platinum  ore,  which 
occurs  as  rounded  grains  or  flattened  scales  in  alluvial  de- 
posits, is  first  digested  with  mercury  or  dilute  aqua  regia 
to  remove  the  gold,  silver,  and  copper;  and  then  with  con- 
centrated aqua  regia,  which  changes  all  the  platinum  and 
a  very  little  iridium  into  soluble  compounds,  leaving  behind 
an  alloy  of  iridium  and  osmium.  From  the  clear  solution 
the  platinum  and  iridium  are  precipitated  by  ammonium 
chloride  as  compounds,  which,  on  heating,  yield  the  metals 
as  a  spongy  mass.  This  spongy  platinum  is  melted  in  a 
lime  crucible  with  an  oxyhydrogen  flame  or  in  the  electric 
furnace,  and  hammered  while  hot  into  sheet  platinum. 

554 


PLATINUM   AND   ASSOCIATED  METALS  555 

The  very  small  amount  of  iridium  is  seldom  removed  from 
the  metallic  platinum. 

Properties  and  Uses  of  Platinum.  —  Platinum  is  a  lustrous, 
gray-white,  soft  metal.  It  is  malleable  and  ductile,  and 
usually  appears  in  commerce  in  the  form  of  wire,  sheet,  and 
dishes.  Sheet  platinum  is  cut  into  squares  —  the  familiar 
platinum  foil  of  the  laboratory,  or  made  into  crucibles, 
dishes,  and  stills.  Its  use  in  these  forms  is  due  partly  to 
its  infusibility  and  partly  to  its  resistance  to  acids  and  other 
corrosive  chemicals.  Although  it  is  attacked  by  fused 
caustic  alkalies,  low  melting  metals,  aqua  regia,  and  a  few 
other  substances,  it  is  practically  indispensable  in  the  chemi- 
cal laboratory  and  is  used  in  many  chemical  processes  in- 
volving accurate  analysis.  Platinum  is  a  good  conductor  of 
electricity,  and  large  quantities  are  consumed  in  electrical 
apparatus,  especially  incandescent  electric  light  bulbs. 
Short  pieces  of  wire  are  fused  through  the  glass  at  the  base  of 
the  bulb  and  thereby  serve  as  electric  conductors.  Platinum 
is  the  only  metal  thus  far  found  which  is  perfectly  adapted 
to  this  use,  because  it  has  the  same  coefficient  of  expansion 
as  glass.  Recently  platinum  has  come  into  use  as  jewelry. 
Platinum  has  a  specific  gravity  of  about  21,  which  is  higher 
than  that  of  any  known  substance,  except  osmium  and  irid- 
ium. The  melting  point  is  about  1755°  C.  In  the  form  of 
a  black,  porous  mass  it  is  called  spongy  platinum,  and  a  still 
finer  form  is  called  platinum  black.  Both  forms  absorb 
large  volumes  of  gases ;  and  if  a  current  of  gas  is  directed 
against  the  metal,  the  gas  often  takes  fire.  Coherent  plati- 
num has  the  same  property  to  a  less  degree,  for  it  becomes 
red-hot  if  held  in  a  stream  of  illuminating  gas,  and  often  ig- 
nites the  gas.  Finely  divided  platinum  is  used  as  the  cata- 
lyzer in  the  contact  method  for  making  sulphuric  acid.  Plati- 
num forms  alloys  with  other  metals,  and  should  never  be 


556  INORGANIC  CHEMISTRY 

heated  with  lead,  similar  metals,  or  their  compounds,  since 
the  alloys  have  a  low  melting  point.  With  iridium,  however, 
it  forms  a  very  hard  alloy  of  which  the-  standard  meters  are 
made. 

Compounds  of  Platinum.  —  Chloroplatinic  acid  (H2PtCl6) 
is  a  brownish,  deliquescent  solid  prepared  by  treating 
platinum  with  aqua  regia  and  evaporating  the  solution  to 
crystallization.  By  carefully  heating  chloroplatinic  acid 
in  a  current  of  chlorine,  platinum  tetrachloride  (PtCLt)  is 
obtained  ;  it  resembles  chloroplatinic  acid  in  color  but  is 
not  deliquescent.  The  acid  forms  salts,  the  best  known 
being  the  yellow,  crystalline  potassium  chloroplatinate 
(K2PtCl6)  and  ammonium  chloroplatinate 


Palladium  is  used  in  chemical  analysis  to  absorb  hydrogen,  in 
making  scientific  instruments,  as  a  catalyst,  and  as  a  substitute  for 
platinum.  A  native  (as  well  as  an  artificial)  alloy  of  iridium  and  os- 
mium, called  iridosmine,  is  used  to  tip  gold  pens.  Iridosmine  is 
often  called  osmiridium. 


PROBLEMS  AND  EXERCISES  (Review) 

1.  (a)  How  much  sodium  chloride  is  needed  to  prepare  a  kilo- 
gram of  hydrogen  chloride  ?     (6)  A  Mloliter  at  20°  C.  and  765  mm.  ? 
(c)  What  weight  of  chlorine  can  be  obtained  from  the  gas  prepared 
in  (6)  ? 

2.  Find  the  weight  of  a  mixture  of  210  cc.  of  oxygen  and  790  cc. 
of  nitrogen.      (Standard  conditions.) 

3.  Discuss  the  following  topics:     (a)  Electrolytes  depress  the 
freezing  point  abnormally  ;   (6)  ions  migrate  to  their  respective  elec- 
trodes ;    (c)  tests  are  for  ions. 

4.  (a)  The  equivalent  weight  of  zinc  was  found  by  experiment  to 
be  32.6  and  the  specific  heat  .095.     What  is  the  approximate  atomic 
weight  ?     (6)  Similarly  for  tin,   equiv.   wt.  =  29.42  and  sp.   ht.  = 
.056.     Calculate  the  approximate  atomic  weight. 


APPENDIX 

i.  The  Metric  System  of  weights  and  measures  is  used  in  physics 
and  chemistry. 

It  is  based  on  the  meter.  This  is  the  unit  of  length,  and  it  is  a  little 
longer  than  a  yard.  Its  exact  length  is  39.37  inches  —  a  number  to 
remember.  Lengths  shorter  than  a  meter  are  called  decimeters, 
centimeters,  and  millimeters.  Deci-  means  .1,  centi-  means  .01,  and 
milli-  means  .001.  Lengths  longer  than  a  meter  are  called  deca- 
meters, hectometers,  and  kilometers.  Deca-  means  10,  hecto-  means 
100,  kilo-  means  1000.  Notice  that  all  these  relations  are  decimal. 
Hence,  a  meter  contains  10  decimeters,  100  centimeters,  or  1000 
millimeters.  It  is  also  evident  that  10  millimeters  equal  1  centi- 
meter, 10  centimeters  equal  1  decimeter, -and  that  1000  meters  equal 
1  kilometer.  The  millimeter,  centimeter,  decimeter,  and  meter  are 
the  denominations  most  frequently  used  in  physical  science  to  express 
length,  though  very  long  distances  are  expressed  in  kilometers.  It 
is  advisable  to  remember  that  — 

1  decimeter  =  about  4  inches. 
30  centimeters  =  about  1  foot. 
2.5  centimeters  —  about  1  inch. 

The  customary  abbreviations  of  the  linear  denominations  are 
meter,  m.;  decimeter,  dm.;  centimeter,  cm.;  and  millimeter,  mm. 

The  unit  of  weight  is  the  gram.  It  is  a  small  weight,  being  only 
about  one  thirtieth  of  an  ounce.  A  five-cent  coin  weighs  approxi- 
mately five  grams.  The  weights  of  small  objects  and  the  small 
quantities  used  in  chemical  analysis  are  expressed  in  terms  of  the 
gram.  The  weights  of  heavy  objects  and  large  quantities  are  often 
expressed  in  terms  of  the  kilogram,  which  is  1000  times  heavier  than 
the  gram.  Just  as  the  meter  is  subdivided,  so  the  gram  is  subdivided 
decimally  into  smaller  weights  called  the  decigram,  centigram,  and 
milligram.  Therefore,  a  gram  contains  10  decigrams,  100  centi- 
grams, or  1000  milligrams;  and  a  kilogram  contains  1000  grams. 

557 


558  INORGANIC  CHEMISTRY 

The  gram,  decigram,  centigram,  milligram,  and  occasionally  the 
kilogram,  are  used  in  physical  science,  though  the  gram  is  the  most 
common  denomination.  For  example,  if  an  object  weighs  2  grams, 
2  centigrams,  and  5  milligrams,  the  weight  is  expressed  as  2.025 
grams;  similarly,  5  milligrams  is  often  expressed  as  .005  gram. 
Preferable  abbreviations  of  the  weight  denominations  are  gram,  gm.; 
decigram,  dg.;  centigram,  eg.;  milligram,  mg.;  and  kilogram,  kg. 

The  unit  of  volume  is  the  liter.  It  is  slightly  larger  than  a  quart, 
and  is  used  for  both  dry  and  liquid  measure.  As  in  the  case  of  the 
meter  and  the  gram,  the  liter  is  subdivided  into  denominations  called 
the  deciliter,  etc.  But  these  fractional  denominations  are  seldom 
used.  That  is,  small  volumes  are  not  expressed  as  decimal  fractions 
of  a  liter,  but  as  cubic  centimeters.  A  liter  contains  1000  cubic 
centimeters,  and  parts  of  a  liter  are  designated  by  the  proper  num- 
ber of  cubic  centimeters.  For  example,  one  half  a  liter  is  called  500 
cubic  centimeters,  one  fourth  is  250  cubic  centimeters,  one  tenth  is 
100  cubic  centimeters;  two  liters  is  often  called  2000  cubic  centi- 
meters. The  relation  of  cubic  centimeters  to  a  liter  is  simple.  The 
French  chemists  who  devised  the  metric  system  first  found  the  length 
of  the  meter  by  measuring  a  part  of  the  meridian  passing  near  Paris. 
Subsequently,  they  constructed  a  vessel  equal  to  the  capacity  of  a 
cubical  vessel  having  edges  10  centimeters  long;  such  a  cubical  vessel 
contains,  of  course,  1000  cubic  centimeters.  The  capacity  of  this 
vessel  they  named  the  liter.  Therefore,  the  liter  and  1000  cubic 
centimeters  are  identical,  whatever  the  substance  measured  —  a  fact 
to  remember.  The  abbreviation  of  liter  is  1.,  and  of  cubic  centimeter 
is  cc.  or  cm.3. 

The  relation  between  meter  and  liter  has  been  shown.  Another 
important  relation  should  be  noted.  The  cubical  vessel  named  the 
liter  may,  of  course,  be  filled  with  any  substance;  if  it  is  filled  with 
pure  water  at  4°  C.,  that  weight  of  water  is  called  a  kilogram.  There- 
fore in  the  case  of  water  the  following  relation  exists:  1  liter,  1  quart, 
and  1000  cubic  centimeters  weigh  approximately  the  same  as  1  kilo- 
gram, 1000  grams,  and  2.2  pounds.  Since  many  liquids  have  about 
the  same  specific  gravity  as  water,  this  general  relation  is  useful,  and 
should  be  learned.  It  is  clear  from  the  relation  just  given  that  1 
cubic  centimeter  of  water  weighs  1  gram  —  a  fact  to  remember,  since 
this  relation  enables  us  to  convert  volume  into  weight,  and  vice  versa. 

The  relation  between  the  units,  multiples,  and  submultiples  of  the 
metric  system  is  shown  in  the  — 


APPENDIX 


559 


TABLE  OF  THE  METRIC  SYSTEM 


LENGTH 

WEIGHT 

VOLUME 

NOTATION 

Kilometer 
Hectometer 
Decameter 
METER 
Decimeter 
Centimeter 
Millimeter 

Kilogram 
Hectogram 
Decagram 
GRAM 
Decigram 
Centigram 
Milligram 

Kiloliter 
Hectoliter 
Decaliter 
LITER 
Deciliter 
Centiliter 
Milliliter 

1000. 
100. 
10. 

1. 

0.1 
0.01 
0.001 

From  this  table  it  is  evident  that  10  milligrams  equal  1  centigram, 
10  centigrams  equal  1  decigram,  10  decigrams  equal  1  gram,  and 
so  on. 

The  relation  of  the  metric  system  to  weights  and  measures  in  com- 
mon use  is  shown  by  the  — 

TABLE  OF  METRIC  EQUIVALENTS 


1  meter 

=  39.37  inches 

linch 

=  2.54    centime- 

ters 

1  kilometer 

=  0.62  mile 

1  mile 

=  1.6  kilometers 

1  centimeter 

=  0.39  inch 

1  cubic  inch 

=  16.39       cubic 

centimeters 

1  liter 

=  0.908  quart  (dry) 

1  quart  (liq.) 

=  0.9465  liter 

1  liter 

=  1.056  quarts  (liq.) 

1  pound  (avoir.) 

=  0.4536       kilo- 

gram 

1  gram 

=  15.452  grains 

.1  ounce  (avoir.) 

=  28.35  grams 

1  kilogram 

=  2.2        pounds 

1  ounce  (troy) 

=  31.1  grams 

(avoir.) 

1  metric  ton 

=  2204  pounds 

1  grain  (apoth.) 

=  0.0648  gram 

The  passage  from  the  English  to  the  metric  system  may  be  accom- 
plished by  utilizing  the  — 

TABLE  OF  METRIC  TRANSFORMATION 


To  CHANGE 

MULTIPLY  BY 

Inches  to  centimeters               .... 

2  54 

Centimeters  to  inches    

0.3937 

Cubic  inches  to  cubic  centimeters 

16387 

Cubic  centimeters  to  cubic  inches    

0061 

Ounces  to  grams  (avoir.)    

28.35 

Grams  to  ounces  (avoir.)    

0.0353 

Grains  to  grams    

0.0648 

Grams  to  grains    . 

15  43 

560 


INORGANIC  CHEMISTRY 


PROBLEMS 

1.  What  is  the  abbreviation  of  gram,  centigram,  liter,  meter,  cubic 
centimeter,  centimeter,  decimeter,  milligram? 

2.  Express  (a)  1  liter  in  cubic  centimeters,  (6)  2  1.  in  cc.,  (c)  1  meter 
in  centimeters,  (d)  250  cm.  in  dm.,  (e)  1  kg.  in  grams,  (/)  250  gm. 
in  mg. 

3.  Add  2  kg.,  5  dg.,  2  eg.,  4  gm.,  and  7  mg.,  and  express  the  sum 
in  grams. 

4.  How  many  cc.  in  a  Hter? 

5.  What  is  the  weight  in  grams  of  (a)  1  liter  of  water,  (6)  250  cc., 
(c)  500  cc.,  (d)  721  cc.? 

6.  Express  in  grams  (a)  721  kg.,  (6)  62  mg.,  (c)  245  eg.,  (d)  84  dg. 

7.  Express  (a)  40  meters  in  inches,  (6)  25  kilograms  in  pounds, 
(c)  54  grams  in  ounces,  (d)  72  grams  in  grains,  (e)  75  liters  in  quarts 
(Hq.). 

2.  The  Thermometer  in  scientific  use  is  the  centigrade.  The  boil- 
ing point  of  water  on  this  thermometer  is  100,  and  the  freezing  point 
is  0  (Fig.  83).  The  equal  spaces  between  these  points  are  called 
degrees.  The  abbreviation  for  centigrade  is  C.,  and  for  degrees  is  °. 
Thus,  the  boiling  point  of  water  is  100°  C.  Degrees  below  zero  are  al- 
ways designated  as  minus,  e.g.  —  12°  C.  means  12  degrees  below  zero. 
The  thermometer  in  popular  use  is  the  Fahrenheit. 
On  this  instrument  the  boiling  point  of  water  is 
212°  and  the  freezing  point  is  32°  above  zero  (Fig. 
100  212  83).  The  abbreviation  for  Fahrenheit  is  F. 

To  change  Fahrenheit  degrees  into  the  equiva- 
lent centigrade  degrees,  subtract  32  and  multiply 
the  remainder  by  f,  or  briefly  — 

C.  =  $  (F. -32). 

To  change  centigrade  degrees  into  the  equivalent 
Fahrenheit  temperature,  multiply  by  f  and  add  32 
to  the  product,  or  briefly  — 

F.  =  f  C.+32. 

The  point  —  273  °C.  is  called  absolute  zero.  Abso- 
lute temperature  is  reckoned  from  this  point. 
Degrees  on  the  absolute  scale  are  found  by  adding 

273  to  the  readings  on  the   centigrade  thermometer.     Thus,  273° 

absolute  is  0°  C.,  274°  absolute  is  +  1°  C.,  etc. 


FIG.  83.  —  Ther- 
mometers. 


APPENDIX 


561 


PROBLEMS 

1.  Change  into  Fahrenheit  readings  the  following  centigrade  read- 
ings:   (a)  60.5,  (6)  40,  (c)  92,  (d)  -5,  (e)  0,  CO  100,  (g)  860,  (/i)  -40. 

2.  Change  into  centigrade  readings  the  following  Fahrenheit  read- 
ings:   (a)  207,  (6)  180,  (c)  0,  (d)  -30,  (e)  212,  (/)  100,  (0)  -40,  (ft)  270. 

3.  Express  the  following  centigrade  readings  in  absolute  readings: 
(a)  0,  (6)  24,  (c)  -13,  (d)  -260. 

3.  Crystallization.  —  Most  substances  in  passing  from  a  liquid  or 
a  gas  into  a  solid  assume  a  definite  shape.  This  change  is  called 
crystallization,  and  the  substances  are  said  to  crystallize  or  to  form 
crystals.  Crystals  are  produced  by  (1)  evaporating  a  solution, 
(2)  cooling  a  melted  solid,  or  (3)  cooling  a  vapor.  Thus,  sodium 
chloride  crystals  are  formed  by  evaporating  a  salt  solution;  sulphur 
crystals,  by  melting  and  then  cooling  sulphur;  and  iodine  crystals, 
by  heating  iodine  in  a  test  tube.  These  methods  are  called,  respec- 
tively, evaporation,  fusion,  and  sublimation. 

As  a  rule,  each  substance  has  an  individual  crystal  form  by  which 
it  can  be  distinguished.  Although  there  are  thousands  of  different 
crystals,  all  belong  to  one  of  six  classes  or  systems.  This  classifica- 
tion is  based  upon  two  assumptions:  (1)  all  crystals  contain  certain 
lines  called  axes,  and  (2)  the  surfaces  or  faces  are  grouped  around  the 
axes  in  definite  positions.  The  axes  connect  angles,  edges,  or  faces, 
which  are  similarly  situated  on  opposite  sides  of  the  crystal.  The 
bounding  planes  or  faces  are  arranged  symmetrically  around  the  axes, 
which  also  determine  (by  their  lengths  and  relative  positions)  the 
positions  of  the  bounding  planes.  For  example,  the  cube  has  three 
equal  axes  at  right  angles  to  one  another  and  terminating  in  the 
center  of  each  of  the  six  bounding  surfaces. 

The  following  is  a  brief  description  of  the  six  systems  of  crystal- 
lization:— 


FIG.  84.  —  Isometric  crystals  (cube,  octahedron,  dodecahedron). 


(1)  Isometric.  —  This  has  three  equal  axes  intersecting  at  right 
angles.     The  simplest  forms  are  the  cube,  octahedron,  and  dodeca- 


562 


INORGANIC  CHEMISTRY 


hedron  (Fig.  84).     Substances  crystallizing  in  this  system  are  dia 
mond,  common  salt,  alum,  fluor  spar,  iron  pyrites,  and  garnet. 


FIG.  85.  —  Tetragonal  crystals. 

(2)  Tetragonal.  —  This  has  three  axes  at  right  angles;  but  one 
axis  is  shorter  or  longer  than  the  other  two,  which  are  equal.  The 
common  forms  are  the  prism,  pyramid,  and  their  combinations 
(Fig.  85).  Tin  dioxide  and  zircon  form  tetragonal  crystals. 


FIG.  86.  —  Orthorhombic  crystals. 

(3)  Orthorhombic.  —  This  has  three  unequal  axes  intersecting  at 
right  angles.  Common  forms  are  the  prism,  pyramid,  and  their  com- 
binations (Fig.  86).  Potassium  nitrate,  barium  sulphate,  topaz,  and 
native  sulphur  crystallize  in  this  system. 


FIG.  87,  —  Hexagonal  crystals. 

(4)  Hexagonal.  —  This  has  four  axes:  three  are  equal  and  inter- 
sect at  60°  in  the  same  plane;  the  fourth  is  longer  or  shorter  than 
the  others  and  is  at  right  angles  to  their  plane.  It  is  a  complex 


APPENDIX 


563 


system.  Common  forms  are  the  prism,  pyramid,  rhombohedron, 
scalenohedron,  and  their  combinations  (Fig.  87).  In  this  extensive 
system  are  found  quartz,  calcite,  beryl,  corundum,  and  ice  (see  Figs. 
60  and  71). 

(5)  Monoclinic.  —  This  has  three  unequal  axes:  two  cut  each  other 
obliquely,  and  the  third  is  at  right  angles  to  the  plane  of  the  other 
two.  Common  forms  are  combinations  of  prisms.  It  is  a  complex 
system,  but  includes  many  substances,  e.g.  sulphur  deposited  by 
fusion,  sodium  carbonate,  borax,  gypsum,  and  ferrous  sulphate 
CFig.  88). 


FIG.  88.  —  Monoclinic  crystal. 


FIG.  89.  —  Triclinic  crystals. . 


(6)  Triclinic.  —  This  has  three  unequal  axes,  all  intersecting  at 
oblique  angles.  Common  forms  are  complex  combinations.  Copper 
sulphate,  potassium  dichromate,  boric  acid,  and  several  minerals 
form  triclinic  crystals  (Fig.  89). 

4.  Vapor  Pressure.  —  The  value  of  a  in  the  formula  in  Chapter  V 
can  be  found  in  the 

TABLE  or  VAPOR  PRESSURE: 


t 

a 

t 

a 

t 

a 

t 

a 

10 

9.18 

16 

13.57 

22 

19.66 

28 

28.10 

.5 

9.49 

.5 

14.00 

.5 

20.26 

.5 

28.93 

11 

9.81 

17 

14.45 

23 

20.88 

29 

29.79 

.5 

10.14 

.5 

14.91 

.5 

21.52 

.5 

30.66 

12 

10.48 

18 

15.38 

24 

22.18 

30 

31.56 

.5 

10.83 

.5 

15.87 

.5 

22.85 

.6 

32.47 

13 

11.19 

19 

16.37 

25 

23.65 

31 

33.41 

.5 

11.56 

.5 

16.88 

.5 

24.26 

.5 

34.37 

14 

11.94 

20 

17.41 

26 

24.99 

32 

35.36 

.5 

12.33 

.5 

17.95 

.5 

25.74 

.5 

36.37 

15 

12.73 

21 

18.50 

27 

26.61 

33 

37.41 

.5 

13.14 

.5 

19.07 

.5 

27.29 

.5 

38.47 

564 


INORGANIC  CHEMISTRY 


5.  Atomic  Weights.  —  The  following  is  a  table  of 

INTERNATIONAL  ATOMIC  WEIGHTS  (1916) 


ELEMENT 

SYM- 
BOL 

AT.  WT. 

ELEMENT 

SYM- 
BOL 

AT.   WT. 

Aluminium      .     . 

Al 

27.1 

Molybdenum  .     . 

Mo 

96.0 

Antimony  .     .     . 

Sb 

120.2 

Neodymium    .     . 

Nd 

144.3 

Argon     .... 

A 

39.88 

Neon      .... 

Ne 

20.2 

Arsenic  .... 

As 

74.96 

Nickel    .... 

Ni 

58.68 

Barium  .... 

Ba 

137.37 

Niton     .... 

Nt 

222.4 

Bismuth      .     .     . 

Bi 

208.0 

Nitrogen     .     .     . 

N 

14.01 

Boron     .... 

B 

11.0 

Osmium      .     .     . 

Os 

190.9 

Bromine      .     .     . 

Br 

79.92 

Oxygen       .     . 

0 

16.00 

Cadmium    .     .     . 

Cd 

112.40 

Palladium  .     .     . 

Pd 

106.7 

Caesium  .... 

Cs 

132.81 

Phosphorus     .     . 

P 

31.04 

Calcium      .     .     . 

Ca 

40.07 

Platinum    .     . 

Pt 

195.2 

Carbon   .... 

C 

12.005 

Potassium  .     . 

K 

39.10 

Cerium   .... 

Ce 

140.25 

Praseodymium    . 

Pr 

140.9 

Chlorine      .     .     . 

Cl 

35.46 

Radium      .     .     . 

Ra 

226.0 

Chromium  .     .     . 

Cr 

52.0 

Rhodium    .     .     . 

Rh 

102.9 

Cobalt    .... 

Co 

58.97 

Rubidium  .     .     . 

Rb 

85.45 

Columbium  *    .     . 

Cb 

93.5 

Ruthenium     .     . 

Ru 

101.7 

Copper  .... 

Cu 

63.57 

Samarium       .     . 

Sa 

150.4 

Dysprosium     .     . 

Dy 

162.5 

Scandium  .     .     . 

Sc 

44.1 

Erbium  .... 

Er 

167.7 

Selenium    .     . 

Se 

79.2 

Europium    .     .     . 

Eu 

152.0 

Silicon  .... 

Si 

28.3 

Fluorine      .     .     . 

F 

19.0 

Silver    .... 

Ag 

107.88 

Gadolinium      .     . 

Gd 

157.3 

Sodium       .     .     . 

Na 

23.00 

Gallium  .... 

Ga 

69.9 

Strontium  .     . 

Sr 

87.63 

Germanium      .     . 

Ge 

72.5 

Sulphur      .     . 

S 

32.06 

Glucinum  2       .     . 

Gl 

9.1 

Tantalum  .     .     . 

Ta 

181.5 

Gold       .... 

Au 

197.2 

Tellurium  .     .     . 

Te 

127.5 

Helium   .... 

He 

4.00 

Terbium     .     .     . 

Tb 

159.2 

Holinium     .     .     . 

Ho 

163.5 

Thallium    .     ,     . 

Tl 

204.0 

Hydrogen    .     .     . 

H 

1.008 

Thorium     .     .     . 

Th 

232.4 

Indium   .... 

In 

114.8 

Thulium     .     .     . 

Tm 

168.5 

Iodine     .... 

I 

126.92 

Tin             .     .     . 

Sn 

118.7 

Iridium  .... 

Ir 

193.1 

Titanium    .     .     . 

Ti 

48.1 

Iron   .... 

Fe 

55.84 

Tungsten    .     .     . 

W 

184.0 

Krypton      .     .     . 

Kr 

82^92 

Uranium    .     .     . 

U 

238.2 

Lanthanum      .     . 

La 

139.0 

Vanadium  .     .     . 

V 

51.0 

Lead  

Pb 

207.20 

Xenon   .... 

Xe 

130.2 

Lithium  .... 

Li 

6.94 

Ytterbium 

Lutecium    .     .     . 

Lu 

175.0 

(Neoytterbium) 

Yb 

173.5 

Magnesium      .     . 

Mg 

24.32 

Yttrium      .     .     . 

Yt 

88.7 

Manganese  . 

Mn 

64.93 

Zinc  

Zn 

65.37 

Mercury      .     .     . 

Hg 

200.6 

Zirconium  .     .     . 

Zr 

90.6 

Or  Niobium,  Nb. 


*  Or  Beryllium,  Be. 


INDEX 


Absolute  alcohol,  314. 

And  centigrade,  43,  44. 

Temperature,  43. 

Zero,  44,  560. 

Acetates,  319,  320,  450,  515. 
Acetic  acid,  317. 

Glacial,  318. 

Ions,  318. 

Test,  320. 
Acetone,  316. 
Acetylene,  290. 

And  calcium  carbide,  285. 

As  illuminant,  291. 

Burner,  292. 

Composition,  291. 

Explosion,  290. 

Flame,  291,  292. 

Generation,  292. 

Thermal  properties,  292. 
Acid,  Acetic,  317. 

Boracic,  374. 

Boric,  374. 

Butyric,  319,  322. 

Carbolic,  329. 

Carbonic,  280. 

Chamber,  348. 

Chloroplatinic,  556. 

Citric,  320. 

Disulphuric,  352. 

Fluosilicic,  384. 

Fuming  nitric,  223. 

Fuming  sulphuric,  351. 

Glacial  acetic,  318. 

Glacial  phosphoric,  395. 

Hydriodic,  189,  190,  371. 

Hydrobromic,  368. 

Hydrochloric,  203,  421. 

Hydrocyanic,  330. 

Hydrofluoric,  364. 

Hydro fluosilicic,  384. 

Hydrosulphuric,  337. 


Acid,  Hypochlorous,  200,  201,  208. 

Lactic,  319,  326,  423. 

Manganic,  520. 

Metaboric,  375. 

Metaphosphoric,  395. 

Metasilicic,  381,  383. 

Muriatic,  203. 

Nitric,  216. 

Nitrous,  221. 

Oleic,  319. 

Orthophosphoric,  395. 

Orthosilicic,  381,  383. 

Oxalic,  319. 

Palmitic,  319,  321,  322,  323. 

Prussic,  330. 

Pyroligneous,  317. 

Pyrophosphoric,  396. 

Pyrosulphuric,  352. 

Silicic,  380,  381,  382. 

Stearic,  319,  321,  322,  323. 

Sulphuric,  345. 

Sulphurous,  342,  343. 

Tartaric,  319. 

Tetraboric,  375. 
Acid  anhydrides,  163. 

Calcium  carbonate,  281. 

Denned,  150. 

Dibasic,  157. 

Monobasic,  157. 

Of  air,  282. 

Oxides,  162. 

Phosphate,  396. 

Potassium  fluoride,  365. 

Potassium  tartrate,  319,  320. 

Reaction,  150. 

Salt,  157,  165. 

Sodium  carbonate,  166,  422. 

Sodium  sulphate,  166. 

Sulphates,  351. 

Tribasic,  158. 
Acidity,  158. 
565 


566 


INDEX 


Acids,  149. 

Acetic  series,  317. 

And  hydrogen  ions,  150. 

And  non-metals,  150. 

And  oxides,  162. 

Chlorine,  160. 

Composition,  150. 

Dissociation,  degree,  164. 

Fatty  series,  317. 

Formulas,  259. 

Hydrogen  ions,  146. 

In  butter,  322. 

Nomenclature,  160. 

Organic,  317,  319. 

Phosphorus,  395. 
Actinium,  531. 
Adsorption,  271. 
Air,  115. 

Acid  of,  282. 

Alkaline,  213. 

And  atmosphere,  115. 

And  carbon  dioxide,  121. 

And  water  vapor,  120. 

Bad,  120. 

Composition,  118. 

Fixed,  282. 

Liquid,  17,  125. 

Mixture,  124. 

Solubility,  65,  125. 

Weight  of  liter,  115. 
Albumin,  397. 
Alchemists,  458. 
Alcohol,  314. 

Absolute,  314. 

Denatured,  314. 

Ethyl,  312,  314. 

Methyl,  314. 

Test,  320. 

Triacid,  321. 

Wood,  314. 
Alcoholic  fermentation,  315. 

Liquors,  315. 
Alcohols,  312. 
Aldehyde,  formic,  316. 
Alizarin,  330. 
Alkali,  151,  422. 

Family,  435. 

Fixed,  161. 

Metals,  417,  435. 

Soda,  422. 
Alkalies,  151,  161. 


Alkalies.     See  Alkali  and  Base. 
Alkaline  air,  213. 

Earth  family,  475. 

Reaction,  151. 
Allotrope,  274. 

Allotropism,  274,  336,  387,  393. 
Alloys,  416. 

Aluminium,  447,  496. 

Antimony,  512. 

Copper,  446. 

Fusible,  407. 

Lead,  512. 

Manganese,  517. 

Platinum,  555. 

Silver,  455. 

Tin,  508. 

Zinc,  484. 

Alpha  particles,  529,  531. 
Alum,  500. 

Cake,  499. 

Chrome,  525. 

Composition,  501. 

History,  501. 

Ions,  501. 

Iron,  547. 
Alumina,  496. 

Aluminates,  494,  496,  497,  498. 
Aluminium,  493. 

Acetate,  319,  501. 

Alloys,  447,  496. 

And  radium,  529,  530. 

Bronze,  447. 

Carbide,  289. 

Chloride,  502. 

Compounds,      hydrolysis,     499, 
501,  502. 

Electrolytic    manufacture,    493, 
494. 

Gems,  497. 

Hydroxide,  497. 

In  crust,  11,  493. 

Ions,  503. 

Metal  and  non-metal,  502. 

Occurrence,  493. 

Oxide,  496. 

Phosphate,  497. 

Properties,  494. 

Reduction  by,  416,  495. 

Silicate,  503. 

Sulphate,  499. 

Test,  502. 


INDEX 


567 


Alumino-thermic  method,  496. 
Aluminum.     See  Aluminium. 
Alundum,  496,  488. 
Amalgam,  416,  488. 

Gold,  488. 

Sodium,  419. 

Tin,  508. 

Zinc,  488. 

Amalgamation,  414,  453. 
Amethyst,  379,  497. 
Ammonia,  210,  213. 

Anhydrous,  212. 

Composition,  216. 

Copper  compounds,  450. 

Dissociation,  212. 

Formation,  210,  219. 

Liquid,  212,  215. 

Muriate,  436. 

Of  commerce,  211. 

Preparation,  210. 

Properties,  211. 

Refrigerant,  215. 

See  Ammonium  hydroxide. 

Soda  process,  422. 

Water,  213. 

Ammoniacal  liquor,  211,  297,  299. 
Ammonium,  214,  438. 

Carbonate,  438. 

Chloride,  435,  436. 

Chloroplatinate,  556. 

Compounds,  213,  214,  435. 

Hydroxide,  211,  213,  214,  215. 

Ion,  438. 

Molybdate,  527. 

Nickelous  sulphate,  551. 

Nitrate,  221,  437. 

Polysulphide,  437. 

Sulphate,  437. 

Sulphide,  437. 

Sulphide,  yellow,  437. 
Ammo  no-compounds,  491. 
Amorphous,  266. 

Carbon,  266. 

Sulphur,  336. 
Amyl  acetate,  321. 

Valerate,  321. 
Analysis,  qualitative,  78,  340. 

Quantitative,  78. 

Spectrum,  438. 

Water,  54. 
An  atmosphere,  116. 


Anglesite,  510. 

Anhydride,  163,  224,  280,  413. 
Anhydrite,  470. 
Anhydrous,  72. 

Copper  sulphate,  72. 
Aniline,  329. 
Animal  charcoal,  325. 
Anion,  denned,  134. 
Anode,  141,  180. 
Anthracene,  330. 
Anthracite  coal,  268. 
Antichlor,  203,  344. 
Antimony,  404. 

Alloys,  512. 

Compounds,  405,  406. 

Metalloid,  410. 

Name,  405. 

Oxides,  405 

Oxychloride,  406. 

Test,  406. 
Apatite,  362,  391. 
Aqua  ammonia,  211,  213. 

Fortis,  219. 

Regia,  224,  460. 
Argol,  319. 
Argon,  122,  441. 

In  air,  118. 

Monatomic,  251. 
Aristotle,  115. 
Arrhenius,  133. 
Arsenic,  401. 

Acids,  403. 

Antidote,  402,  546. 

Marsh's  test,  404. 

Metalloid,  410. 

Oxide,  402. 

Poisoning,  402. 

Properties,  402. 

Pyrites,  401. 

Salts,  403. 

Sulphides,  403. 

Sulpho-salts,  404. 

Test,  404. 

Trichloride,  402. 

Trioxide,  402. 

White,  402. 
Arsenious  oxide,  402. 
Arsine,  404. 
Artificial  diamonds,  264, 

Graphite,  266. 

Stone,  383.  ' 


568 


INDEX 


Asbestos,  479. 
Ash,  coal,  268,  270. 

Potassium  compounds,  434. 

Silica,  380. 
Assaying,  453,  454. 
Assimilation  of  nitrogen,  131. 
Atmosphere,  115. 

An,  116. 

Ingredients,  116. 

Pressure,  116. 

See  Air. 
Atomic  groups,  ammonium,  214. 

Combinations,  253,  257. 

Displacement,  257. 

Ions,  134. 

Valence,  255. 

Atomic  theory,  91,  227,  233,  234. 
Atomic  weights,  96. 

And    equivalent    weights,    231, 
244. 

And  properties,  357. 

And  specific  heat,  246. 

And  symbols,  97. 

And  valence,  254. 

Approximate,  245. 

Carbon,  243. 

Determination,  228,  242-248. 

Exact,  246,  247. 

International  table,  564. 

Oxygen,  243. 

Standard,  228,  237. 

Zinc,  247,  248. 
Atoms,  92. 

And  ions,  134. 

And  molecules,  92,  95. 

Combination,  253,  257. 

Decay,  530,  531. 

Decomposition,  95,  530,  531. 

Displacement,  258. 
Atoms  in  molecule,  250. 

Argon,  251. 

Arsenic,  402. 

Arsenious  oxide,  403. 

Bromine,  367. 

Cadmium,  251,  486. 

Chlorine,  251. 

Fluorine,  363. 

Hydrofluoric  acid,  365. 

Hydrogen,  236,  251. 

Iodine,  370. 

Mercury,  251,  488. 


Atoms.  Nitrogen,  251. 

Oxygen,  237,  251. 

Ozone,  251. 

Phosphorus,  251,  394. 

Potassium,  251. 

Sodium,  251,  419. 

Sulphur,  251,  335. 

Zinc,  251,  483. 
Attraction  and  repulsion,  141. 
Auric  chloride,  461. 
Avogadro's  hypothesis,  232-235. 
Azote,  117. 
Azurite,  442,  449. 

Babbitt's  metal,  484. 

Baking  powder,  320,  423,  438. 

Soda,  423. 
Balard,  368. 
Barite,  474. 
Barium,  473,  474,  475,  476. 

Ions,  475. 

Oxides,  17,  19,  101,  187,  474. 

Oxides,  equilibrium,  187. 

Sulphate,  equilibrium,  194. 
Barometer,  116. 
Baryta  water,  474. 
Barytes,  474. 
Base,  152. 

And  hydroxyl  ions,  152. 

And  metals,  151. 

And  oxides,  162. 

Composition,  151. 

Diacid,  158. 

Dissociation,  164. 

lonization,  164. 

Monacid,  158. 

Nomenclature,  161. 

See  Alkali. 

Triacid,  158 
Basic,  151. 

Anhydrides,  163. 

Lining,  540. 

Oxides,  163. 

Reaction,  158. 

Salt,  158,  165. 

See  Alkaline. 
Basicity,  157. 
Basil  Valentine,  345. 
Bath  metal,  446. 
Battery,  electric,  178. 
Bauxite,  498. 


INDEX 


569 


Becher,  25. 

Beckmann  apparatus,  240. 

Becquerel,  528. 

Beer,  315. 

Beet  sugar,  325. 

Residues,  428,  433. 
Bell  metal,  447. 
Benzaldehyde,  316. 
Benzene,  299,  329. 
Benzine,  294. 
Benzol,  329. 
Bergman,  282. 
Berlin  blue,  549. 
Berthollet,  90. 
Beryl,  383,  497. 
Berzelius,  85. 
Bessemer  steel,  539. 
Beta  particles,  529. 
Bichromates,  523. 
Binary  compounds,  162. 
Biscuit  ware,  503. 
Bismite,  406. 
Bismuth,  406. 

Oxides,  407. 

Oxychloride,  407. 

Subnitrate,  407. 

Test,  407. 
Bismuthinite,  406. 
Bisulphite,  calcium,  328. 

Soda,  344. 
Bitter  almonds,  316. 
Bituminous  coal,  268. 
Bivalent  elements,  254,  256,  257. 
Black,  282,  481. 
Black  ash,  421. 

Damp,  289. 

Lead,  265,  510. 
Blast  furnace,  534. 

Lamp,  38. 
Bleaching,  chlorine,  200,  202. 

Hydrogen  dioxide,  87. 

Sodium  dioxide,  427. 

Sulphur  dioxide,  342. 
Bleaching  powder,  201. 
Bleach  liquors,  209. 
Blooms,  538. 
Blowpipe,  flame,  309. 

Mouth,  308. 

Oxyhydrogen,  37. 
Bluestone,  448. 
Blue  vitriol,  448. 


Boiler  scale,  473. 

Boiling  point,  elevation,  74,  140. 

Water,  60,  560. 
Bone  ash,  391,  453. 

Cupel,  454. 
Bone  black,  272. 
Bones,  391,  396. 
Boracic  acid,  374. 
Boracite,  374. 
Borax,  375. 

Bead,  375,  376. 

Hydrolysis,  376. 

Reaction,  376. 

Use,  376,  473. 
Bordeaux  mixture,  449. 
Boric  acid,  374. 
Boride,  carbon,  374 
Bornite,  442. 
Boron,  374. 

Nitride,  374. 

Oxide,  374,  376. 

Test,  375. 
Bort,  264. 
Boussingault,  119. 
Boyle's  law,  46,  233. 
Brand,  391. 
Brandy,  315. 
Brass,  446. 
Braunite,  517. 
Bread,  328,  423. 
Breathing,  26. 
Brimstone,  334. 
Brin's  process,  17. 
Britannia  metal,  508. 
Bromides,  365,  366,  367,  368. 
Bromine,  365. 

Discovery,  368. 

Formula,  367. 

Preparation,  366. 

Properties,  367. 

Water,  367. 
Bronze,  446. 

Aluminium,  447. 

Phosphor,  446. 

Silicon,  447. 
Brucite,  480. 
Bunsen,  air,  118. 

Burner,  305. 

Flame,  305. 

Spectroscope,  417,  440. 
Burette,  156. 


570 


INDEX 


Burner,  acetylene,  292. 

Bunsen,  305. 

Self-lighting,  34. 
Burning,  23. 

See  Combustion. 
Butter,  322. 
Butyric  acid,  322. 

Cadmium,  486. 

Hydroxide,  486. 

Ions,  486. 

Sulphide,  486. 

Test,  486. 

Caesium,  417,  435,  440. 
Calamine,  482. 
Calcite,  464. 
Calcium,  463. 

Acid  carbonate,  467. 

Acid  phosphate,  396. 

Bicarbonate,  467. 

Bisulphite,  328. 

Borate,  375. 

Carbide,  284,  285,  290,  464-467. 

Carbonate,  464,  488. 

Carbonate,  acid  and  normal,  281. 

Chlorate,  432. 

Chloride,  463,  471. 

Compounds,    and    hard    water, 
473. 

Cyanamide,  472. 

Fluoride,  362,  363,  364,  365. 

Hydride,  464. 

Hydroxide,  66,  469. 

Hypochlorite,  201. 

In  crust,  11,  463. 

Ions,  472. 

Light,  38. 

Magnesium  carbonate,  478,  481. 

Nitrate,  217. 

Nitride,  464. 

Occurrence,  11,  463. 

Oxalate,  472. 

Oxide,  467. 

Phosphate,  391-394,  396,  400. 

Preparation,  463. 

Properties,  463,  464. 

Sulphate,  470. 

Sulphide,  421,  472. 

Sulphite,  acid,  328. 

Test,  472. 
Caliche,  427. 


Calomel,  489. 
Calorie,  155,  175,  530. 
Calorific  value,  coal,  269. 
Calorimeter,  175. 
Candle  power,  300,  309. 
Cane  sugar,  324. 

See  Sugar. 
Cannizzaro,  235. 
Caramel,  324. 
Carat,  diamond,  265. 

Gold,  460. 
Carbide,  aluminium,  289. 

Calcium,  284. 
Carbides,  275,  284. 

And  petroleum,  294. 
Carbohydrates,  324. 
Carbolic  acid,  329. 
Carbon,  263-275. 

Amorphous,  263,  266,  274. 

As  fuel,  275. 

Atomic  weight,  243. 

Bisulphide,  352. 

Boride,  374. 

Chemical  properties,  274. 

Disulphide,  65,  352,  353. 

Gas,  273,  297. 

Monoxide,  282,  283,  300,  536. 

Oxides,  275. 

See  Carbon  dioxide. 

Silicide,  385. 

Test,  272. 

Tetrachloride,  313. 
Carbona,  313. 
Carbonado,  264. 
Carbonate,  280. 

Acid,  281. 

Normal,  281. 
Carbon  dioxide,  276. 

And  combustion,  276. 

Baking  powder,  423. 

Composition,  281. 

Critical  temperature,  278. 

Formation,  276. 

History,  282. 

In  atmosphere,  121. 

Liquid,  278. 

Mines,  277. 

Occurrence,  276. 

Preparation,  277. 

Relation  to  life,  279. 

Solid,  278. 


INDEX 


571 


Carbon  dioxide,  solubility,  277,  278. 

Test,  276,  470. 
Carbonic  acid,  280. 

Anhydride,  280. 

Ions,  280. 
Carborundum,  385,  496. 

Furnace,  386. 
Carboxyl,  317. 

Carnallite,  366,  428,  429,  478. 
Carnotite,  476. 
Casein,  326. 
Cassiterite,  505. 
Cast  iron,  536,  537. 
Castner  method,  sodium,  418. 
Catalysis,  185,  288,  313. 

Chlorine,  197. 

Sulphuric  acid,  349. 

Water,  200. 
Cathode,  141,  180. 
Cations,  134. 
Caustic  alkali,  151. 

Lime,  468. 

Lunar,  456. 

Potash,  433. 

Soda,  424. 

Cavendish,  25,  36,  39,  82,  118. 
Caves,  465. 
Celestite,  474. 
Cell,  electrolytic,  140,  179. 

Semi-permeable,  137. 

Storage,  180. 

Voltaic,  178. 
Celluloid,  329. 
Cellulose,  328. 

Nitrates,  328. 
Cement,  469,  480. 

Plaster,  471. 

Cementation  process,  542. 
Cementite,  537. 
Centigrade  and  absolute  temperature, 

43,  44. 

Cerium,  515,  516. 
Cerussite,  510. 
Chalcedony,  379. 
Chalcocite,  442. 
Chalcopyrite,  442,  532. 
Chalk,  466. 
Chalybeate  water,  53. 
Chamber  acid,  348. 
Change,  kinds,  102. 

Physico-chemical,  5. 


Change.    See  Chemical  and  Physical 

Action. 
Chaptal,  129. 
Charcoal,  271. 

Animal,  272,  325. 

Distillation,  272. 

Filter,  271. 

Kiln,  271. 

Pit,  271. 

Wood,  271,  272,  273. 
Charles'  law,  43,  233. 
Chemical  action,  4,  17,  102. 

And  electricity,  177. 

And  heat,  171. 

And  hydrogen,  32. 

And  light,  169. 

And  oxygen,  17,  22. 

And  solution,  76,  184. 

Kinds,  102. 

Chemical  compounds,  13,  78. 
Chemical  energy,  7,  169,  177. 

See  Energy. 
Chemical  equation,  101,  102,  107. 

See  Equation. 
Chemical  equilibrium,  184,  186-195. 

See  Equilibrium. 
Chemical  equivalents,  229. 

See  Equivalents. 
Chemical  properties,  5. 
Chemical  reaction,  6. 
Chemical  symbols,  10,  11,  12,  13,  97, 
134,  564. 

And  atoms,  97. 

And  ions,  134. 
Chemistry,  defined,  1. 

Organic,  310. 
Chile  saltpeter,  426. 
China,  503. 
Chinese  white,  484. 
Chlorate  of  potash,  432. 
Chloride,    199,    203,    206,    207,    224, 
335. 

Of  lime,  201. 

Chlorination  process,  459. 
Chlorine,  196-201. 

Acids,  160,  208. 

And  bromides,  366. 

And  iodides,  369. 

And  water,  79. 

Available,  201. 

Hydrate,  199. 


572 


INDEX 


Chlorine,  Ionic,  145,  208. 

Liquid,  199. 

Oxides,  209. 

Preparation,  196,  197,  424,  425. 

Properties,  198. 

Replacing  power,  367,  369. 

Water,  79,  198. 
Chloroform,  313. 
Chlorophyll,  279,  532. 
Chloroplatinic  acid,  556. 
Choke  damp,  289. 
Chromates,  523,  527. 
Chrome,  alum,  525. 

Iron  ore,  522. 

Yellow,  515,  525. 
Chromic  acid,  527. 

Compounds,  526. 

Hydroxides,  526. 

Oxide,  526. 

Sulphate,  526. 
Chromite,  522,  527. 
Chromites,  526,  527. 
Chromium,  522. 

Alum,  525. 

As  metal  and  non-metal,  525. 

Compounds,  523,  524,  525,  526. 

Ions,  527. 

Mordant,  525. 

Oxides,  526,  527. 

Preparation,  495,  522. 

Test,  524,  525. 

Valence,  526,  527. 
Chromous  compounds,  526. 
Cinnabar,  486,  487,  489. 
Citric  acid,  320. 
Classification  of  Elements,  355. 

Periodic,  357,  361. 

Table,  358. 
Clay,  503. 

Ironstone,  548. 
Coal,  266. 

Anthracite,  268,  269. 

Beds,  267. 

Bituminous,  268,  269,  295. 

Brown,  269. 

Calorific  value,  269. 

Classes,  268. 

Composition,  268. 

Distillation,  295. 

Distribution,  270. 

Fossil,  267. 


Coal,  Gas,  295. 

Gas,  composition,  300. 

Gas,  plant,  295,  296. 

Hard,  269. 

Map,  270. 

Products  from,  297. 

Section,  267. 

Soft,  269. 

Tar,  297. 
Cobalt,  552. 

Ions,  146. 

Test,  553. 
Cobaltite,  552. 
Coins,  bronze,  446. 

Gold,  461. 

Nickel,  447,  550. 

Silver,  455. 
Coke,  273,  385. 
Colemanite,  374,  375. 
Collodion,  328. 
Colloid,  381. 
Colloidal  solution,  382. 

Arsenic  trisulphide,  404. 

Ferric  hydroxide,  546. 

Gold,  461. 

Suspension,  382. 
Combination,  19,  22. 

And  equation,  102. 
Combining  capacity,  253,  257. 

Of  atomic  groups,  257. 

Of  atoms,  257. 
Combustion,  24,  25. 

And  heat,  171. 

And  hydrogen,  37. 

And  light,  170. 

Broad  use,  35. 

Old  theory,  25. 

Ordinary,  276. 

Present  explanation,  25. 

Products,  304. 

Spontaneous,  24. 

Common  salt.     See  Sodium  chloride. 
Composition,  air,  118. 

Ammonia  gas,  216. 

Carbon  dioxide,  281. 

Coal,  268. 

Earth's  crust,  10,  11. 

Illuminating  gas,  300. 

Of  a  compound,  78. 

Organic  compounds,  310. 

Percentage,  100, 


INDEX 


573 


Composition,  Sulphur  dioxide,  342. 

Water,  78-85. 
Compounds,  13,  78. 

Formula,  98,  248. 

Concentration  and  equilibrium,  189- 
195. 

Fraction,  194. 

Ionic,  193,  195. 

Maximum,  67. 

Of  ore,  414. 

Of  solutions,  63. 
Concrete,  469. 
Condenser,  water,  55. 

Illuminating  gas,  295. 
Conductivity,  electrical,  132,  133. 

See  Electrolysis. 
Condy's  liquid,  520. 
Conservation,  energy,  8,  169. 

Matter,  7,  169. 

Constitution,     organic     compounds, 
311. 

Formula,  312. 

Contact  method,  sulphuric  acid,  348. 
Converter,  539. 
Cooking  soda,  423. 
Copper,  442. 

Acetate,  450. 

Alloys,  446. 

Ammonia  compounds,  450. 

And  nitric  acid,  219. 

And  sulphuric  acid,  341. 

Blister,  444. 

Carbonates,  442,  449. 

Chlorides,  447. 

Coins,  446. 

Compounds,  447. 

Displacement,  445,  450,  451. 

Electrolytic,  444. 

Family,  461. 

Fluoride,  364. 

Glance,  442. 

Ions,  146,  147,  450. 

Matte,  443. 

Metallurgy,  443. 

Nitrate,  449. 

Ore  map,  533. 

Oxide,  442,  447. 

Poisonous,  447,  448. 

Properties,  445. 

Purification,  445.    • 

Pyrites,  442. 


Copper,  Replacement,  445,  450,  451. 

See  Cupric  and  Cuprous. 

Sulphate,  143,  444,  447,  448,  449. 

Sulphate,   hydrolysis,    166,    167, 
448. 

Sulphide,  442,  449. 

Test,  450. 

Uses,  445. 
Copperas,  547. 
Coquina,  466. 
Coral,  467. 
Cordite,  323. 
Corpuscles,  96. 
Corrosive  sublimate,  489. 
Corundum,  496,  497. 
Courtois,  370. 
Cracking  petroleum,  300. 
Cream  of  tartar,  320,  423,  433. 
Critical  pressure,  128. 

Acetylene,  290. 

Ammonia,  128. 

Carbon  dioxide,  278. 

Ethylene,  289. 

Hydrogen,  128. 

Oxygen,  128. 
Critical  temperature,  128,  290. 

Air,  128. 

Ammonia,  212. 

Carbon  dioxide,  128,  278. 

Ethylene,  289. 

Nitrous  oxide,  222. 

Oxygen,  128. 

Sulphur  dioxide,  128,  342. 
Crockery,  504. 
Crocoisite,  522. 
Crocoite,  522. 
Crocus,  546. 

Crucible  process,  steel,  542. 
Cryolite,  362,  389,  493,  501. 
Crystallization,  69,  561. 

And  solution,  70. 

Water  of,  71,  72. 
Crystalloids,  382. 
Crystals,  69,  73. 

Systems,  561. 
Gullet,  387. 
Cupellation,  453. 
Cupric  compounds,  447. 

Ferrocyanide,  450. 

Ions,  447,  448. 

Oxide,  448, 


574 


INDEX 


Cupric  compounds.     See  Copper. 

Sulphate,  448. 
Cuprite,  448. 
Cuprous  compounds,  447. 

Ions,  447. 

Oxide,  448. 

See  Copper. 
Curie,  528. 
Curve,  solubility,  68. 
Cyanamide,  472. 
Cyanide,  mercury,  330. 

Potassium,  434. 

Potassium  auri-,  461. 

Potassium  silver,  455. 

Process,  415,  459. 

Sodium,  427. 
Cyanides,  330. 

Iron,  548. 
Cyanogen,  330. 

Ions,  146. 
Cymogene,  294. 

Dalton,  atomic  theory,  92. 

Multiple  proportions,  90. 
Davy,   alkali   metals,    81,    177,    198, 
222,  264,  308,  370,  417,  418, 
419,  463. 

Deacon  process,  197. 
Decay,  27,  276,  288. 

Animal  matter,  210,  216. 
Decomposition,  18. 

And  equations,  102. 

Double,  103,  154,  207. 

Spontaneous,  530. 
Decrepitation,  430. 
Definite  proportions,  89. 
Deflagration,  220. 
Degree  of  dissociation,  135,  163. 
Dehydrated  compounds,  72. 
Deliquescence,  74,  424,  426,  432,  433, 

480. 

Denatured  alcohol,  314. 
Density,  19. 

Gases,  48. 

Oxygen,  19. 

Water,  56. 
Depression    of    freezing     point,    74, 

139. 

Destructive  distillation,  289. 
Determination,  atomic  weights,  228, 
242-248. 


Determination,   Equivalent  weights, 
230. 

Molecular     weights,     238,     240, 

241,  242. 

Dewar,  34,  125,  128. 
Dew-point,  120. 
Dextrose,  324,  325,  326. 
Diamond,  263. 

Allotrope,  274. 

And  radium,  476. 

Artificial,  264. 

Cheap,  380. 

Crystals,  264. 

Cullinan,  265. 

Gems,  264,  265. 

Royal,  265. 
Diaspore,  498. 
Diastase,  315,  327. 
Diatomaceous  earth,  379. 
Dibasic  acid,  157. 
Dicalcium  phosphate,  401. 
Dichromates,  523,  527. 
Diffusion,  34,  122. 
Disilicate,  383. 
Disinfectant,  316,  330,  520. 
Disodium  phosphate,  396. 
Displacement,  downward,  198. 

Of  equilibrium,  191-195. 

Of  metals,  450. 

Upward,  211. 
Dissociation,  acids,  163. 

Ammonia,  212. 

Bases,  163. 

Degree,  135. 

Electrolytic,  133. 

In  solution,  134,  163. 

Salts,  163. 

Table,  164. 

Water,  61,  62. 
Distillation,  54. 

Coal,  295. 

Destructive,  289. 

Dry,  210,  317. 

Petroleum,  294. 

Water,  54. 

Wood,  272,  317. 
Disulphuric  acid,  352. 
Dolomite,  478,  481. 
Double  decomposition,  103,  207. 

Neutralization,  154. 
Double  salt,  428,  501. 


INDEX 


575 


Downward  displacement,  198. 
Drinking  water,  53. 

And  ozone,  29. 
Dulong  and  Petit,  245. 
Dumas,  84,  119,  264,  356. 
Dutch  metal,  446. 

Process,  white  lead,  514. 
Dyads,  257. 
Dyeing,  501. 
Dynamite,  323. 

Earth's  crust,  11. 
Effervescence,  278. 
Efflorescence,  72,  422,  426. 
Electric  charges,  134,  142. 

Current,  179. 

Furnace,  173,  285,  385. 
Electrical  energy,  measurement,  183. 
Electricity  and  chemical  change,  177. 

And  solution,  132. 
Electrochemical  equivalent,  183. 

Series,  metals,  451. 
Electrochemistry,  182. 
Electrodes,  141,  179,  266. 
Electrolysis,  179,  183. 

Aluminium,  493. 

Applications,  181. 

Calcium  chloride,  463. 

Carnallite,  478. 

Copper  sulphate,  143,  444. 

Extraction  by,  415. 

Fused  salts,  144,  417,  478,  493. 

Gold,  460. 

Hydrochloric  acid,  141. 

Illustrations,  143. 

Interpretation,  180. 

Rusting,  544. 

Silver,  455. 

Sodium  chloride,  424. 

Sodium  hydroxide,  417. 

Sodium  sulphate,  143. 

Solution,  140. 

Theory,  140-144,  180. 

Water,  80,  143. 
Electrolytes,  133,  136. 
Electrolytic  cell,  140,  179,  455. 
Electrolytic  dissociation,  133,  544. 
Electrolytic  process,  chlorine,  198. 

Sodium  hydroxide,  417. 
Electrolytic  solution,  133. 

Boiling  point,  140. 


Chemical  behavior,  144. 

Freezing  point,  139. 
Electron,  96,  529,  531. 
Electro-negative  ions,  134,  146,  147. 
Electroplating,  182. 
Electro-positive  ions,  134,  146,  147. 
Electrosilicon,  380. 
Electrotyping,  181,  182. 
Elements,  8-13,  358,  564. 

Acid-forming,  150,  413. 

Base-forming,  151,  413. 

Bivalent,  254,  256,  257. 

Classification,  355. 

Decay,  530,  531. 

Families,  356,  359. 

Groups,  356,  359. 

In  crust,  11. 

Molecule  of,  95,  250. 

Numerical  relations,  356. 

Periodic  classification,  357. 

Periodic  table,  358. 

Properties,  acid,  150,  355,  413. 

Properties,  basic,  151,  355,  413. 

Quadrivalent,  254,  255,  257. 

Quinquivalent,  257. 

Radioactive,  476,  528-531. 

Tables,  10,  11,  12,  564. 

Trivalent,  254,  255,  257. 

Univalent,  254-257. 

Valence,  254-257. 
Elevation  of  boiling  point,  74,  140. 
Emerald,  497. 
Emery,  496. 
Empirical  formula,  312. 
Endothermic   compounds,    176,   292, 

352. 
Energy,  7,  8,  274. 

And  light,  170. 

Chemical,  7,  169,  177. 

Conservation,  8. 

Electrical,  177,  183. 

Heat,  measurement,  175. 

See  Electrolysis. 

Storage  cell,  180. 
Enzymes,  315,  325,  327,  328. 
Epsom  salts,  480. 
Equation,  101. 

And  calculations,  111. 

Gas,  252. 

Ionic,  154. 

Making,  107. 


576 


INDEX 


Equation,  Molecular,  251. 

Neutralization,  153,  156. 

Ordinary,  102. 

Preliminary,  18. 

Reversible,  189. 

Thermal.  176. 

Volume,  252. 
Equilibrium,  61,  70,  184,  186-195. 

Ammonia,  212. 

And  solution,  189,  193-195. 

Barium  oxides,  187. 

Barium  sulphate,  194. 

Concentration,  191-195. 

Constant,  191. 

Displacement,  191-195. 

Heat,  191,  192,  224. 

Hydriodic  acid,  189,  190,  371. 

Hydrochloric  acid,  204. 

Hydrogen,  187-191. 

Le  Chatelier's  law,  191,  224. 

Nitric  acid,  217. 

Nitrogen  oxides,  224. 

Phosphorus  pentachloride,  398. 

Saturated  solution,  70. 

Steam  and  iron,  187,  188,  191. 

Water  vapor,  61.        • 
Equivalent  weights,  229. 

And  atomic  weights,  231,  244. 

And  valence,  254. 

Determination,  230. 

Table,  230. 
Equivalents,  chemical,  183. 

Electrochemical,  183. 

See  Equivalent  weights. 
Erosion,  51. 
Esters,  320. 
Etching,  365. 
Ether,  ethyl,  316. 

Solubility  in  water,  66. 

Sulphuric,  316. 
Ethyl,  311,  312. 

Acetate,  320. 

Alcohol,  312,  314. 

Butyrate,  321. 

Ether,  316. 
Ethylene,  289. 

Evaporation  of  water,  57-61. 
Exothermic  compounds,  176. 
Explosions,  acetylene,  290. 

Coal  mine,  289. 

Hydrogen,  35. 


Factors,  106. 

Fahrenheit  thermometer,  560. 

Faraday,  127,  179,  183,  199. 

Law,  183. 
Fats,  321,  322. 
Fehling's  solution,  326. 
Feldspar,  503. 
Ferment,  314,  318. 
Fermentation,  315. 

Acetic,  318. 

Alcoholic,  277,  315. 

Bread,  327. 

Sugar,  325,  326. 
Ferric  compounds,  544. 

Chloride,  548. 

Ferrocyanide,  549. 

Hydroxide,  546. 

Oxide,  546. 

Sulphate,  547. 

Sulphide,  547. 

Sulphocyanate,  549. 
Ferricyanides,  548,  549. 
Ferrite,  546. 
Ferrochrome,  543. 
Ferrocyanides,  450,  548,  549. 
Ferromanganese,  517,  540. 
Ferrosilicon,  543. 
Ferroso-ferric  oxide,  546. 
Ferro tungsten,  543. 
Ferrous  compounds,  544. 

Carbonate,  548. 

Chloride,  548. 

Chromite,  522. 

Ferric  oxide,  546. 

Ferricyanide,  549. 

Ferrocyanide,  549. 

Hydroxide,  546. 

Oxide,  546. 

Sulphate,  222,  546. 

Sulphide,  547. 
Fertilizer,    131,   400,   401,   426,   434, 

437,  541. 

Filter,  charcoal,  271. 
Fire,  24. 

And  carbon  oxides,  282. 

Damp,  288. 

Extinguisher,  278. 

See  Combustion. 
Fireworks,  24,  474. 
Fixation,  nitrogen,  131. 
Fixed  air,  282. 


INDEX 


577 


Fixed  alkali,  161. 
Flame,  301. 

Acetylene,  290,  291. 

And  gauze,  307. 

Bunsen,  305,  306,  307. 

Candle,  302,  303. 

Hydrogen,  36. 

Lamp,  302. 

Luminosity,  304. 

Luminous,  302,  305. 

Non-luminous,  305,  306. 

Ordinary  gas,  204. 

Oxidizing,  308. 

Oxy acetylene,  27. 

Oxy  hydro  gen,  37. 

Parts,  302,  303. 

Structure,  302,  303. 

Reducing,  308. 
Flashing  point,  294. 
Flint,  379. 

Flowers  of  sulphur,  334. 
Fluid  magnesia,  481. 
Fluorides,  365. 
Fluorine,  362. 

Compounds,  362. 

Formula,  363. 

Isolation,  362. 
Fluorite,  362. 
Fluor  spar,  362. 
Fluosilicic  acid,  384. 
Flux,  415,  535. 
Fool's  gold,  547. 
Formaldehyde,  316. 
Formalin,  316. 
Formula,  97. 

And  valence,  258. 

Calculation,  101. 

Constitutional,  260. 

Determination,  248. 

Empirical,  312. 

Graphic,  260,  312,  317. 

Molecular,  248. 

Rational,  312. 

Simplest,  249,  250. 

Structural,  260,  312,  317. 

Weight,  238. 

Writing,  258. 

Forward  reaction,  188,  190,  192. 
Franklinite,  482. 
Freezing  point,  560. 

And  molecular  weight,  239. 


Freezing  point,  And  solution,  139. 

Depression,  139. 

Determination,  240,  241,  248. 

Solution.  74, 
Fructose,  326. 
Fruit  sugar,  326. 
Fuming  acid,  nitric,  223. 

Sulphuric,  351. 
Furnace,  blast,  534. 

Electric,  173,  174. 

Open  hearth,  541. 

Reverberatory,  415. 
Fusible  alloys,  407. 

Metals,  407,  486. 

Gahnite,  482. 

Galena,  510,  514. 

Galvanized  iron,  483. 

Gangue,  414. 

Garnet,  497. 

Garnierite,  550. 

Gas,  burner,  self-lighting,  34. 

Carbon,  273,  297. 

Coal,  295. 

Equation,  252. 

Flame,  302,  303. 

Illuminating,  295. 

Liquor,  211. 

Marsh,  288. 

Natural,  288,  295. 

Olefiant,  289. 

Pintsch,  299. 

Producer,  32,  283. 

Reaction,  189. 

Water,  299. 
Gases,  adsorption  by  charcoal,  271. 

And  "pressure,  41,  42,  46. 

And  temperature,  41,  42-46. 

Density,  48. 

Gay-Lussac's  law,  224,  233. 

Inert,  123. 

Kinetic  theory,  232. 

Liquefaction,  127. 

Molecular  formula,  250. 

Solution,  64. 
Gasolene,  294. 
Gastric  juice,  204. 
Gay-Lussac,  83,  346,  370. 

Law,  224,  233. 

Tower,  346. 
Gems,  aluminium,  497. 


578 


INDEX 


Gems,  Artificial,  497. 
German  silver,  447. 
Geyserite,  384. 
Glacial  acetic  acid,  318. 

Phosphoric  acid,  395. 
Glass,  387. 

Annealing,  388. 

Blowing,  389. 

Colored,  389. 

Crown,  389. 

Cut,  388. 

Etching,  365. 

Flint,  388. 

Hard,  389. 

Kinds,  388. 

Lead,  388. 

Manufacture,  388. 

Quartz,  380. 

Water,  383. 

Window,  389. 
Glauber  salt,  426. 
Glazing,  503,  504. 
Globigerina  ooze,  466. 
Glover  tower,  346. 
Glucose,  326. 
Glycerin,  321,  322. 
Glycerol,  323. 
Glyceryl,  321. 

Oleate,  321. 

Palmitate,  321. 

Stearate,  321. 
Gold,  457. 

Alloys,  458,  461. 

Amalgam,  488. 

Carat,  460. 

Coin,  461. 

Colloidal,  461. 

Compounds,  461. 

Cyanide  process,  459. 

Electrolytic  purification,  460. 

Fool's,  547. 

History,  458. 

Leaf,  460. 

Metallurgy,  458. 

Mining,  458. 

Occurrence,  458. 

Ores,  458. 

Plating,  461. 

Purification,  459. 

Purity,  458,  460. 

Red,  461. 


Gold,  Separation  from  silver,  459. 

Test,  461. 

Uses,  461. 

White,  461. 

Goldschmidt  method,  496,  497. 
Graham,  34,  382. 
Gram,  557. 

Gram-molecular  volume,  238. 
Gram-molecular    weight,    138,    189, 

240. 

Grape  sugar,  326. 
Graphic  formula,  260,  291,  312. 
Graphite,  265,  266. 

Allotrope,  274. 

Artificial,  266. 
Gravimetric  composition,  air,  118. 

Water,  83. 
Green  fire,  474. 

Vitriol,  546. 
Greenockite,  486. 
Guano,  400. 
Guignet's  green,  526. 
Gun  cotton,  328. 

Metal,  447. 

Gunpowder,  75,  271,  431. 
Gypsum,  470. 

Haemoglobin,  283,  532. 
Halides,  362. 
Halogens,  362. 

Acids,  ions,  372. 

Family,  372. 
Hardness,  water,  473. 
Hausmannite,  517. 
Heat  and  chemical  change,  171. 

Combustion,  176. 

Decomposition,  176. 

Electric  furnace,  173. 

Equilibrium,  191,  192,  224. 

Formation,  176. 

Le  Chatelier's  law,  191,  224. 

Measurement,  175. 

Of  neutralization,  155. 

Of  solution,  75. 

Sources,  173. 

Velocity  reaction,  191,  224. 
Helium,  123. 

And  radium,  123,  531. 

Detection,  441. 

Spectrum,  124. 
Hematite,  532,  533,  534. 


INDEX 


579 


Henry's  law,  65. 
Heptads,  257. 
Heptavalent  elements,  257. 
Heterogeneous  mixture,  189. 
Hexads,  257. 

Hexavalent  elements,  257. 
Homogeneous  mixture,  189,  387. 
Humboldt,  83. 
Humidity,  120. 
Hydrargyllite,  498. 
Hydrate,  71. 

Chlorine,  199. 
Hydraulic  main,  296. 
Hydriodic  acid,  equilibrium,  189,  190, 

371. 

Hydrobromic  acid,  368. 
Hydrocarbons,  288,  299,  313. 
Hydrochloric  acid,  203,  205,  421. 

Commercial,  204. 

Electrolysis,  141. 

Equilibrium,  204. 

Ionic  test,  145. 

Ions,  206. 

Hydrocyanic  acid,  330. 
Hydrofluoric  acid,  364,  379. 

Electrolysis,  363. 
Hydrofluosilicic  acid,  384. 
Hydrogen,  30-39. 

And  acids,  31,  150,  452. 

And  alpha  particles,  529. 

And  chlorine,  35,  199,  205. 

And  iodine,  189,  190. 

And  metals,  452. 

And  oxygen  explosion,  35. 

And  water,  31,  33,  78. 

Arsenide,  404. 

As  metal,  452. 

Atom,  236,  529. 

Bromide,  368. 

Chloride,  203,  236. 

Diffusion,  34. 

Dioxide,  86,  186,  248. 

Displacement,  452. 

Equilibrium,  187-191. 

Explosion,  33,  35. 

Flame,  36,  175. 

Fluoride,  364. 

From  nitric  acid,  219. 

Ions,  146,  150,  452. 

Lavoisier's  experiment,  187. 

Liquid,  128. 


Hydrogen,  Manufacture,  425. 

Molecule,  236. 

Molecular  equation,  252. 

Peroxide.     See  Dioxide. 

Solid,  128. 

Standard,  228,  236. 

Weight  of  liter,  38. 
Hydrogen  sulphide,  337,  339. 

Composition,  339. 

Ions,  339. 

Test,  340. 
Hydrolysis,  166. 

Acid  sodium  sulphate,  166. 

Aluminium  compounds,  499,  501, 
502. 

Antimony  trichloride,  406. 

Bismuth  trichloride,  407. 

Copper  sulphate,  165,  448. 

Ferric  chloride,  166,  548. 

Magnesium  chloride,  480. 

Phosphates,  397. 

Potassium  carbonate,  166. 

Potassium  cyanide,  166. 

Soap,  324. 

Sodium  carbonate,  165,  422. 

Sugar,  325,  326,  327. 

Sulphites,  345. 

Zinc  chloride,  485. 
Hydroquinone,  457. 
Hydroxyl,  147,  152,  154. 
Hydroxyl  ions,  bases,  152. 
Hypo,  352,  457. 
Hypochlorous  acid,  201,  208. 
Hypophosphites,  397. 
Hyposulphite,  sodium,  352. 
Hypothesis,  88. 

Avogadro's,  232-235. 

Ice,  57. 

Plant,  215. 
Iceland  spar,  464. 
Illuminants,  299,  300. 
Illuminating  gas,  295,  299. 

Candle  power,  300,  301. 

Characteristics,  300. 

Composition,  300. 

Impurities,  300. 

Luminosity,  300. 

Tarnishing  by,  454. 
Inert  gases,  123,  124. 
Infusorial  earth,  379. 


580 


INDEX 


Ingots,  542. 

Ink,  456,  547. 

Inorganic  compounds,  310. 

Insecticides,  337,  403. 

Insoluble  substances,  63. 

Invertase,  325. 

Iodides,  368,  369,  371,  373. 

Iodine,  189,  190,  368. 

And  chlorine,  371. 

And  hydrogen,  189,  190. 

And  starch,  185,  327. 

Preparation,  369. 

Properties,  370. 

Purification,  369,  370. 

Saltpeter,  368,  369,  427. 

Seaweed,  368,  369. 

Solution,  372. 

Test,  370. 

Vapor  density,  370. 
lodoform,  312,  372. 
Ionic  equation,  154. 

Concentration,  193,  195. 

Hydrolysis,  167. 

Neutralization,  195. 

Test,  sulphate,  351. 
lonization,  134,  193-195. 

And  salts,  165,  193-195. 

See  Dissociation,  Electrolytic. 
Ions,  defined,  134. 

And  atoms,  134. 

And  electrolysis,  180. 

And  equilibrium,  187,  188,  191, 
193-195. 

And  neutralization,  154. 

And  water,  165,  194,  422,  448, 
449. 

Color,  146. 

Common,  146. 

Complex,  146. 

Effect  of  removal,  191. 

Interaction,  144. 

Migration,   141,   146,   180,  424, 
425. 

Symbols,  134. 

Table,  147. 

Water,  165,  194,  422,  448,  449. 
Iridium,  554,  555,  556. 
Iridosmine,  556. 
Iron,  532. 

Acetate,  319. 

Alum,  547. 


Iron,  And  steam,  187,  188,  190. 
Cast,  536,  537. 
Chemistry  of  smelting,  536. 
Chlorides,  548. 
Compounds,  545. 
Cyanides,  548. 
Bisulphide,  547. 
Electrolytic  preparation,  544. 
Equilibrium,  187,  188,  190. 
Ferrite,  546. 
Flux,  535. 
Galvanized,  483. 
Hydroxides,  545. 
Impurities,  536,  537. 
In  crust,  11,  532. 
Ions,  550. 
Liquor,  319. 
Magnetic  oxide,  546. 
Metallurgy,  534. 
Ore,  532,  533. 
Ore,  chrome,  522. 
Ore,  map,  533. 
Oxides,  297,  545. 
Passive,  544. 
Pig,  536. 
Pure,  544. 
Pyrites,  532,  547. 
Russia,  546. 
Smelting,  534. 
Spiegel,  517. 
Steel,  538. 
Sulphate,  546. 
Sulphides,  547. 
Terms,  532. 
Test,  549. 
Valence,  545. 
Varieties,  536. 
Wrought,  537. 

Jasper,  379. 
Javelle's  water,  209. 

Kainite,  428,  434,  478. 

Kaolin,  383,  503. 

Kerosene,  294. 

Kieserite,  478,  480. 

Kindling  temperature,  25,  172,  30 

307. 

Kinetic  theory,  232. 
Kirchhoff,  440. 
Krypton,  123. 


INDEX 


581 


Labarraque's  solution,  209. 

Lactic  acid,  423. 

Lactose,  326. 

Lake,  501. 

Lampblack,  274. 

Laughing  gas,  222. 

Lavoisier,  25,  27,  31,  37,  39,  83,  117, 

163,  264,  282. 
Law,  88. 

And  theory,  88. 

Boyle,  46,  233. 

Charles,  43,  44,  45,  233. 

Conservation  of  energy,  8. 

Conservation  of  matter,  7,  93. 

Definite  proportions,  89,  94. 

Dulong  and  Petit,  245. 

Faraday,  183. 

Gay-Lussac,  224,  233. 

Henry,  65. 

Le  Chatelier,  191,  224. 

Mass  action,  189-195. 

Multiple  proportions,  90,  94. 

Periodic,  361. 

Specific  heat,  245. 
Lead,  510. 

Acetate,  319,  515. 

Alloys,  512. 

Argentiferous,  453. 

Arsenate,  403. 

Black,  265,  510. 

Carbonate,  510,  513. 

Chloride,  515. 

Chromate,  515,  522,  525. 

Cupellation,  453. 

Dioxide,  513. 

Displacement,  511. 

Fluosilicate,  511. 

Glass,  388. 

In  drinking  water,  512. 

Ions,  515. 

Metallurgy,  510. 

Monoxide,  513. 

Nitrate,  515. 

Oxides,  513. 

Pencils,  266. 

Peroxide,  513. 

Phosphate,  510. 

Poisoning,  512. 

Properties,  511. 

Purification,  511. 

Red,  513. 


Lead,  Storage  cell,  180. 

Sugar  of,  319. 

Sulphate,  510,  515. 

Sulphide,  510,  514. 

Test,  514,  515. 

Tetroxide,  513. 

White,  513,  514. 
Leblanc  process,  421,  433. 
Le  Chatelier's  law,  191,  224. 
Leguminous  plants  and  nitrogen,  131. 
Levulose,  326. 
Liebig,  173,  368. 
Life  and  nitrogen,  130. 

And  oxygen,  26. 

And  phosphorus,  400. 

And  potassium,  434. 
Light  and  chemical  change,  169. 

And  energy,  170. 

And  silver  salts,  456. 
Lignite,  268,  269. 
Lime,  467. 

And  bleaching  powder,  201. 

Caustic,  468. 

Chloride  of,  201. 

Light,  38,  467. 

Milk  of,  470. 

Preparation,  468. 

Quick,  468. 

See  Calcium  oxide. 

Slaking,  171,  467. 

Soda,  469. 

Sulphur,  337. 

Superphosphate,  400. 

Uses,  468. 
Limekiln,  468,  469. 
Limestone,  464,  468. 
Limewater,  469. 

See  Calcium  hydroxide. 
Limonite,  532,  533. 
Link,  valence,  312. 

Fusible,  406. 

Liquefaction  of  gases,  127. 
Liquefied  ammonia,  212,  215. 
Liquid  air,  17,  125,  127. 

Acetylene,  290. 

Ammonia,  212,  215. 

Carbon  dioxide,  278. 

Chlorine,  199. 

Ethylene,  289. 

Hydrogen,  34,  128. 

Oxygen,  20,  126. 


582 


INDEX 


Liquid,  Ozone,  28, 

Sulphur  dioxide,  342s 
Liquids,  solubility,  65, 
Liquor,  alcoholic,  315, 

Distilled,  315. 

Iron,  319, 

Red,  319,  501. 
Litharge,  453,  454,  513* 
Lithium,  417,  435, 
Lithophone,  474, 
Litmus,  acids,  15ft. 

Bases,  151, 

Neutral,  152, 

Salts,  152,  165, 
Loadstone,  546, 
Lockyer,  123, 
Lubricating  oils,  294. 
Luminosity,  300,  302. 
Luminous  paint,  474. 
Lunar  caustic,  456. 
Luster,  412. 
Lye,  151. 

Magnalium,  496. 
Magnesia,  480,  481. 

Alba,  481,  518. 

Black,  518. 

Fluid,  481. 

Mixture,  480. 

See  Magnesium. 

Stone,  518. 
Magnesite,  480,  481. 
Magnesium,  478. 

Alloy,  496. 

Alumina te,  497. 

Ammonium  phosphate,  481. 

Basic  carbonate,  481. 

Bromide,  365,  366,  367. 

Calcium  carbonate,  478,  481. 

Carbonate,  481. 

Chloride,  480. 

Citrate,  481. 

Compounds  and  water,  473. 

Electrolytic  preparation,  478. 

Hydroxide,  480. 

Ions,  481. 

Nitride,  129,  479. 

Oxide,  480. 

Phosphate,  481. 

Properties,  479. 

Stassfurt  salts,  478. 


Magnesium,  Sulphate,  426,  480, 

Test,  481. 

Magnetic  oxide  of  iron,  546. 
Magnetite,  532,  533. 
Malachite,  442,  449, 
Malt,  315. 
Maltose,  315,  327, 
Manganates,  520, 
Manganese,  517. 

Alloys,  517. 

Black  oxide,  518, 

Compounds,  518,  519,  520, 

Dioxide,  196,  517,  518, 

Ferro-,  517. 

Ions,  520. 

Oxides,  517. 

Test,  518,  520. 

Valence,  520. 

Weldon  process,  197,  518. 
Manganic  acid,  520. 
Manganite,  517. 
Manganous  chloride,  518,  520. 

Hydroxide,  518. 

Sulphate,  520. 

Sulphide,  520. 

Mantle,  Welsbach,  309,  516. 
Marble,  464. 
Marchand  tube,  83. 
Marsh  gas,  288. 
Marsh's  test  for  arsenic,  404. 
Mass  action,  189-195. 
Massicot,  513. 
Matches,  399. 
Matter,  conservation,  7. 

Denned,  6. 

Distribution,  11. 

Properties,  1. 
Mendelejeff,  357,  361. 
Mercuric  compounds,  488. 

Chloride,  489,  509. 

Cyanide,  330. 

Ions,  147,  490. 

Nitrate,  490. 

Oxide,  15,  18,  23,  191,  488. 

Sulphide,  486,  487,  489. 
Mercurous  compounds,  488. 

Chloride,  489,  509. 

Ions,  147,  490. 

Nitrate,  490. 
Mercury,  486. 

Amalgams,  488. 


INDEX 


583 


Mercury,  Ammo  no  compounds,  491. 

Compounds,  488. 

Fulminate,  352. 

Preparation,  487. 

Properties,  487. 

See  Mercuric  and  Mercurous. 

Specific  heat,  245. 

Test,  490. 
Metaboric  acid,  375. 

Phosphoric  acid,  395. 

Silicate,  383. 

Silicic  acid,  381,  383. 

Stannic  acid,  507. 

Metal    and     non-metal,     355,    410, 
411. 

Babbitt's,  484. 

Bath,  446. 

Bell,  447. 

Britannia,  508. 

Dutch,  446. 

Gun,  447. 

Hypothetical,  417. 

Muntz,  446. 

Newton's,  407. 

Noble,  224,  452,  458,  554. 

Rose's,  407. 

Speculum,  447. 

Type,  512. 

White,  508. 

Wood's,  407,  486. 
Metallic  ions,  134,  135,  147,  413. 

See  Ions  and  Electrolysis. 
Metalloids,  410. 
Metallurgy,  414. 

See  Individual  metals. 
Metals,  alkali,  417,  435. 

Alkaline  earth,  463,  475. 

Ancient,  413. 

And  bases,  151,  413. 

And  non-metals,  355,  410,  411. 

And  steel,  543. 

Atomic  state,  451. 

Chemical  properties,  413. 

Classification,  411. 

Displacement,  450. 

Electrochemical  series,  451. 

Electromotive  series,  451. 

Electropositive,  451,  452. 

Fusible,  407,  486. 

Ionic  state,  451. 

Ions,  134,  135,  147,  413. 


Metals,  Latin  names,  13. 

Native,  413. 

Occurrence,  413. 

Physical  properties,  412. 

Platinum,  554,  556. 

Preliminary  treatment,  414. 

Preparation,  414. 

Table,  411. 
Metathesis,  103. 
Meteorites,  532,  550. 
Meter,  557. 
Methane,  288,  308. 

In  natural  gas,  295. 

Series,  293. 
Methyl,  311,  312. 

Alcohol,  240,  314. 

Salicylate,  321. 
Metric  system,  557. 

Equivalents,  559. 

Table,  559. 

Transformations,  559. 
Mexican  onyx,  465. 
Meyer,  Victor,  239. 
Microcosmic  salt,  396. 
Migration  of  ions,  142,  180. 

See  Electrolysis. 
Milk  of  lime,  470. 

Sour,  423. 

Sulphur,  337. 
Mineral,  defined,  413. 

Orange,  513. 

Springs,  52. 

Water,  52. 
Minium,  513. 
Mispickel,  401. 
Mixture,  14,  90. 

Air,  124. 

Bordeaux,  449. 

Heterogeneous,  189. 

Homogeneous,  189,  387. 
Moissan,  264,  285,  363,  364,  463. 
Molar  weight,  238. 
Molasses,  325. 
Mole,  138,  176,  189,  240. 
Molecular  depression,  240. 

Equation,  251. 

Formula,  248,  250. 
Molecular  weight,  99,  236,  250. 

And  vapor  density,  236. 

Approximate,  241. 

Depression,  240. 


584 


INDEX 


Molecular     weight,     Determination, 
238,  240,  241,  242. 

Exact,  241. 

Freezing  point,  239. 

Gram,  138,  189,  240. 

Hydrogen  dioxide,  248. 

Hydrogen  standard,  236. 

Oxygen  standard,  236. 

Preliminary,  99. 
Molecules,  92,  95. 

And  atoms,  95. 

And  dissociation,  133,  135. 

And  equations,  251,  252. 

And  equilibrium,   186-189,   192, 
212. 

And  formulas,  97,  248-252. 

And  kinetic  theory,  233. 

See  Atoms  in  molecule. 

Undissociated,  135. 
Molybdenum,  527. 
Monacid  base,  158. 
Monads,  257. 
Monazite,  516. 
Monde  process,  551. 
Monel  metal,  551. 
Monobasic  acid,  157. 
Monoclinic  crystal,  sulphur,  336. 
Mordant,  501,  509,  525,  547. 
Morley,  84. 
Mortar,  470. 
Moth  balls,  329. 
Multiple  proportions,  90,  91. 
Muntz  metal,  446. 
Muriate  of  ammonia,  436. 
Muriatic  acid,  203. 

Names,  acids,  160. 

Bases,  161. 

Salts,  161. 
Naphtha,  294. 
Naphthalene,  329. 
Natural  gas,  295. 
Natural  groups,  356,  357,  358,  359, 

360. 

Natural  waters,  52. 
Negative  electrode,  141,  180. 

Photographic,  457. 
Neon,  123. 

Neutral  reaction,  152. 
Neutralization,  153. 

And  ionization,  154,  194. 


Neutralization,  Equation,  177. 

Heat  of,  155. 
Newton's  metal,  407. 
Niccolite,  550. 

Nicholson  and  Carlisle,  81,  177. 
Nickel,  550. 

Ammonium  sulphate,  550,  551. 

Carbonyl,  551. 

Chloride,  551. 

Coins,  447,  550. 

Compounds,  551. 

Hydroxide,  551. 

Ions,  146,  552. 

Plating,  550. 

Steel,  551. 

Sulphate,  551. 

Sulphide,  551. 

Test,  551. 
Nickeloid,  551. 
Niton,  358,  359,  530,  531. 
Nitrates,  219. 

Behavior  with  heat,  220. 

Formation,  216. 

Test,  220. 
Nitric  acid,  216. 

And  iron,  544. 

And  metals,  219. 

Equilibrium,  217. 

Formation,  216. 

Fuming,  223. 

Ions,  218. 

Manufacture,  218. 

Preparation,  217. 

Properties,  218. 

Test,  220. 
Nitric  oxide,  222. 
Nitride,  magnesium,  129,  479. 
Nitrification,  216. 
Nitrites,  220,  221. 
Nitrobenzene,  329. 
Nitrogen,  129,  131. 

Additional  properties,  129. 

And  life,  130. 

Assimilation,  131. 

Dioxide,  223,  224. 

Family,  408. 

Fertilizer,  437. 

Fixation,  131. 

General  properties,  117.. 

In  atmosphere,  117. 

Liquid,  129. 


INDEX 


585 


Nitrogen,  Oxides,  219,  221. 

Pentoxide,  224. 

Peroxide,  223,  224. 

Tetroxide,  224. 

Trioxide,  224. 
Nitroglycerin,  323. 
Nitrosylsulphuric  acid,  346. 
Nitrous  acid,  221. 
Nitrous  oxide,  221,  438. 
Noble  metal,  452,  458,  554. 
Nomenclature,  161. 

Acids,  160. 

Bases,  161. 

Ions,  134,  146. 

Salts,  157,  161,  162.. 
Non-electrolytes,  133,  136. 
Non-electrolytic  solution,  143. 

Freezing  point,  139. 
Non-luminous  flame,  305. 
Non-metallic  ions,  134,  146,  147. 

See  Ions. 
Non-metals,  150,  355,  410,  411. 

And  acids,  150. 

And  metals,  355,  410. 

Table,  411. 

Nordhausen  sulphuric  acid,  351. 
Normal  salts,  157,  165. 

Solution,  163. 

Occlusion,  34. 
Ocean  water,  53. 

Chlorine  in,  196. 

Composition,  11. 

Solids,  53. 

Table,  11. 
Oil,  lubricating,  294. 

Natural,  322. 

Of  vitriol,  345. 
Olefiant  gas,  289. 
Olein,  321,  322. 
Oleomargarine,  322. 
Olivine,  383. 
Onyx,  379,  465. 
Opal,  379. 
Open-hearth  furnace,  541. 

Process,  541. 
Orange  mineral,  513. 
Ordinary  chemical  equation,  102. 
Ore,  denned,  413. 

Concentration,  414. 

Dressing,  414. 


Ore,  Smelting,  414. 
Organic  acids,  319. 

Compounds,  310,  311,  312. 
Orpiment,  401,  403. 
Orthoclase,  383. 
Orthophosphoric  acid,  395. 
Orthorhombic  crystals,  sulphur,  336. 
Orthosilicic  acid,  381,  383. 
Osmiridium,  556. 
Osmium,  556. 
Osmotic  pressure,  136. 
Oxalic  acid,  319. 

And  carbon  monoxide,  283. 
Oxidation,  22. 

And  decay,  26. 

And  life,  26. 

And  ozone,  28. 

And  reduction,  39. 

Broad  meaning,  490,  509,  545. 

Ionic  standpoint,  545. 

Nitric  acid,  218. 

Potassium  dichromate,  524. 

Potassium  permanganate,  519. 
Oxide,  denned,  22. 

Carbonic,  280. 
Oxides  and  acids,  162,  413. 

And  bases,  162,  413. 

Chlorine,  209. 

Names,  24. 

Nitrogen,  table,  221. 
Oxidizing  agent,  24. 

Hydrogen  dioxide,  87. 

Nitrate,  220. 

Nitric  acid,  218. 
•  Potassium  chlorate,  24. 

Potassium  dichromate,  524. 

Potassium  permanganate,  519. 
Oxidizing  flame,  308. 
Oxone,  17,  427. 
Oxygen,  15-22,  26,  27. 

And  acids,  27,  160. ' 

And  blood,  26. 

And  combustion,  24. 

And  ozone,  28,  29. 

And  water,  79. 

Breathing  pure,  26. 

Brin's  process,  17. 

Discovery,  27. 

Free,  208. 

In  atmosphere,  117. 

Liquid,  20.     See  Liquid  air. 


586 


INDEX 


Oxygen,  Molecule,  237. 

Preparation,  16,  17,  427. 

Properties,  19,  20. 

Relation  to  life,  26. 

Solid,  20. 

Solubility,  65. 

Standard,  228,  236. 

Weight  of  liter,  19. 
Oxy  hydro  gen  blowpipe,  37. 
Oxymuriate  of  tin,  509. 
Ozone,  28,  29. 

Formula,  251. 

Paint,  474,  484,  514. 

Pakfong,  447. 

Palladium,  34,  556. 

Palmitic  acid,  321. 

Palmitin,  322. 

Palm  oil,  322. 

Paper,  328. 

Paracelsus,  39. 

Paraffin,  294. 

Paris  green,  319,  403,  450. 

Paris,  plaster  of,  471. 

Parkes  process,  453. 

Partial  pressure,  60,  65. 

Passive  iron,  544. 

Paste,  glass,  388. 

Pearlash,  433. 

Pentad,  257. 

Percentage  composition,  100. 

Periclase,  480. 

Periodic  classification,  361. 

Law,  361. 

Permanent  hardness,  473. 
Peroxide,  hydrogen,  86,  186,  249. 

Sodium,  427. 
Petrified  wood,  379. 
Petroleum,  293. 

Cracking,  300. 

Distillation,  294. 

Products,  294. 

Refining,  294. 

Well,  293. 
Pewter,  508. 
Phenol,  329. 
Phenyl,  311. 

Philosopher's  stone,  458. 
Phlogiston,  25,  27. 
Phosgene,  284. 
Phosphates,  395,  396. 


Phosphates,  Acid  calcium,  396. 

Dicalcium,  401. 

Disodium,  396. 

Ions,  397. 

Meta-,  396. 

Primary,  396. 

Rock,  400. 

Secondary,  396. 

Slag,  400. 

Sodium  ammonium,  396,  541. 

Tertiary,  396. 

Tests,  396. 

Tricalcium,  391-394,  396. 
Phosphine,  397. 
Phosphonium  compounds,  398. 
Phosphor  bronze,  446. 
Phosphorescence,  476. 
Phosphoric  acids,  391,  392,  395. 

Glacial,  395. 

Ions,  397. 

Phosphoric  oxides,  394. 
Phosphorite,  391. 
Phosphorous  oxide,  394. 
Phosphorus,  391. 

Acids,  391,  392,  395. 

And  nitrogen,  129. 

And  ozone,  28. 

Black,  394. 

Bronze,  446. 

Burns,  393. 

Electrolytic  manufacture,  392. 

Formula,  394. 

Oxides,  394. 

Pentachloride,  398,  399. 

Preparation,  391,  392. 

Properties,  393. 

Red,  394. 

Relation  to  life,  400. 

Sulphide,  399. 

Test,  397. 

Trichloride,  398. 

Yellow,  393. 
Photography,    170,    328,    352,    368, 

456. 

Phylloxera,  337. 
Physical  change,  3. 

Properties,  4. 

Physico-chemical  change,  5. 
Picromerite,  428. 
Pig  iron,  536. 
Pinchbeck,  446. 


INDEX 


587 


Pintsch  gas,  299. 
Pitchblende,  476,  528. 
Plaster,  470. 

Of  Paris,  471. 
Platina,  554. 
Platinum,  554. 

Alloys,  555,  556. 

Arsenide,  554. 

Black,  555. 

Catalyzer,  348,  349,  555. 

Chloroplatinic  acid,  556. 

Foil,  555. 

Metals,  554,  556. 

Properties,  555. 

Sponge,  554,  555. 

Sulphuric  acid  manufacture,  348. 

Tetrachloride,  556. 

Using,  555. 
Plumbago,  265. 
Polariscope,  326. 
Polonium,  531. 
Polyhalite,  428. 
Polysulphides,  437. 
Porcelain,  503. 
Portland  cement,  469. 
Positive  electrode,  141,  180. 
Potash,  432,  433. 

Caustic,  433. 

Prussiate  of,  548. 
Potassium,  428. 

And  water,  429. 

Antimonyl  tartrate,  406. 

Auricyanide,  461. 

Bicarbonate,  432. 

Bichromate,  523,  524. 

Bromide,  368. 

Carbonate,  432,  433. 

Chlorate,  16,  18,  429,  430,  431. 

Chloride,  432. 

Chloroplatinate,  556. 

Chromate,  523. 

Cobaltinitrite,  553. 

Cyanide,  330,  434. 

Bichromate,  523,  524. 

Discovery,  417. 

Electrolytic  manufacture,  432. 

Ferricyanide,  548. 

Ferrocyanide,  548. 

Fluoride,  acid,  365. 

Hydroxide,  433. 

Hypochlorite,  209, 


Potassium,  Iodide,  369. 

Ions,  434. 

Manganate,  520. 

Name,  433. 

Nitrate,  217,  430,  431. 

Nitrite,  430. 

Oxide,  433. 

Perchlorate,  432. 

Permanganate,  518. 

Preparation,  429. 

Properties,  429. 

Relation  to  life,  433,  434. 

Salts,  Stassfurt,  428. 

Silicate,  383. 

Sulphate,  434. 

Sulphocyanate,  330. 

Tartrate,  acid,  319,  320. 

Test,  429. 

Thiocyanate,  330. 
Pottery,  503. 
Powder,  gun,  431. 

Smokeless,  328. 
Pressure,  critical,  128. 

Normal,  41. 

Osmotic,  136. 

Partial,  60,  65. 

Priestley,  27,  82,  118,  173,  213,  222. 
Problems  based  on  equations,  111. 
Producer  gas,  32,  283. 
Properties,  1. 

And  atomic  weights,  357. 

Chemical,  5. 

Classification,  3. 

Physical,  4. 

Radioactive,  476. 
Proteid,  330. 
Protein,  218,  330. 
Proust,  90. 
Prussian  blue,  549. 
Prussiate  of  potash,  548. 
Prussic  acid,  330. 
Puddling,  538. 
Pulmotor,  26. 
Purple  of  Cassius,  461. 
Putty,  467. 
Pyrene,  313. 
Pyrite,  547. 
Pyroligneous  acid,  317. 
Pyrolusite,  517. 
Pyromorphite,  510. 
Pyrophosphates,  397, 


588 


INDEX 


Pyrophosphoric  acid,  396. 
Pyrosulphates,  351. 
Pyrosulphuric  acid,  352. 
Pyrrhotite,  532. 

Quadrivalent  elements,  254,  256,  257. 
Qualitative  analysis,  78,  340. 
Quantitative  analysis,  78,  246,  247. 
Quartation,  459. 
Quartz,  379. 
Quicklime,  468. 
Quicksilver,  487. 
Quinquivalent  elements,  257. 

Radiations,  528,  529. 
Radical,  defined,  154. 

Ammonium,  214. 

Organic,  311,  312. 

Valence,  255,  256. 
Radioactivity,  476,  516,  528-531. 
Radium,  476,  528-531. 

And  helium,  123,  530. 

And  niton,  530. 

Decay,  528. 

See  Radioactivity. 
Ramsay,  118,  122,  123. 
Raoult,  240. 
Rational  formula,  312. 
Rayleigh,  122. 
Rays,  gamma,  529. 

X,  529. 
Reaction,  defined,  101. 

Acid,  150. 

Alkaline,  151. 

Basic,  151. 

Forward,  188,  190,  192. 

Gas,  189. 

Neutral,  152. 

Reverse,  190,  191. 

Reversible,  186. 

Velocity,  184,  185,  188-191. 
Reagents,  6. 
Realgar,  401,  403. 
Red  fire,  474. 

Lead,  513. 

Liquor,  319,  501. 
Reducing  agent,  38. 

Carbon  monoxide,  284. 

Charcoal,  271. 

Hydrogen,  38. 

Stannous  chloride,  490,  509. 


Reducing  flame,  308. 
Reduction,  defined,  38. 

Broad  meaning,  490,  509,  545. 

By  charcoal,  271. 

Ionic  standpoint,  545. 

Ore,  415. 

Refining  petroleum,  294. 
Regenerative  process,  542. 
Relative  humidity,  120. 
Repulsion  and  attraction,  141. 
Respiration,  26. 
Reverberatory  furnace,  415. 
Reverse  reaction,  190,  191. 
Reversible  equations,  189. 
Reversible  reactions,  186-195. 
Reversion,  fertilizer,  401. 
Rhigolene,  294. 
Rhodium,  554. 
Rhodocroisite,  517. 
Richards,  246. 
Rochelle  salt,  320,  448. 
Roll  sulphur,  334. 
Rose's  metal,  407. 
Rouge,  546. 

Rubidium,  417,  435,  440. 
Ruby,  497. 
Rum,  315. 
Rusting,  544. 
Ruthenium,  554. 
Rutherford,  nitrogen,  117. 

Radium,  528. 

Saccharose,  324. 
Safety  lamp,  308. 
Sal  ammoniac,  436. 

Soda,  422. 
Saleratus,  423. 
Salt,  defined,  157,  165. 

Acid,  157,  165. 

Basic,  158,  165. 

Cake,  421. 

Common.     See  Sodium  chloride. 

Double,  428,  501. 

Glauber's,  426. 

Microcosmic,  396. 

Normal,  157,  165. 

Pink,  509. 

Rochelle,  320,  448. 

Tartar,  433. 

Tin,  508. 
Saltpeter,  Chile,  368,  369,  370,  426. 


INDEX 


589 


Salts,  152-159,  165-167,  194,  195. 

Action  on  litmus,  152,  165. 

And  ionization,  165. 

Classification,  156. 

Composition,  153. 

Dissociation,  table,  164. 

Double,  428,  501. 

Epsom,  480. 

In  sea  water,  196. 

Nomenclature,  157,  161,  162. 

Preparation,  158. 

Properties,  152. 

Smelling,  438. 

Solubility  product,  194,  195. 

Stassfurt,  428. 
Sand,  379,  385. 

Blast,  380. 

See  Quartz. 
Saponification,  323. 
Sapphire,  497. 
Satin  spar,  470. 
Saturated  solution,  67. 

And  equilibrium,  194,  195. 
Scheele,  25,  27,  196. 

Green,  403. 
Scrubber,  297. 
Sea  water,  elements  in,  11.     ' 

Solids  in,  53. 

Seaweed  ash,  368,  369,  370. 
Seidlitz  powders,  320,  423. 
Selenite,  470. 
Selenium  glass,  389. 
Semi-permeable  membrane,  137. 
Serpentine,  383. 
Siderite,  532,  533,  548. 
Siemens-Martin  process,  541. 
Silica,  378. 

And  plants,  380. 

See  Quartz. 

Vessels,  380. 
Silicates,  380,  381,  382,  383. 

Alkaline,  383,  384. 

Soluble,  379. 
Siliceous  sinter,  384. 
Silicic  acids,  380,  381,  382. 
Silicides,  385. 
Silicified  wood,  379. 
Silicon,  378. 

Bronze,  447. 

Dioxide,  365,  378. 

Test,  384. 


Silicon,  Tetrafluoride,  365,  384. 
Siloxicon,  384. 
Silver,  452. 

Alloys,  452,  453,  455. 

Bromide,  456. 

Chloride,  452,  456. 

Cleaning,  454. 

Coins,  455. 

Cyanide,  455. 

Determination    atomic    weight, 
247. 

German,  447. 

Horn,  196,  452. 

Iodide,  456. 

Ions,  146,  147,  456. 

Metallurgy,  453. 

Mirror,  455. 

Nitrate,  455. 

Oxidized,  454. 

Phosphates,  397. 

Plating,  455. 

Properties,  454. 

Separation  from  gold,  453,  459. 

Sterling,  455. 

Sulphides,  452. 

Tarnishing,  339,  454. 

Test,  456. 

Simplest  formula,  249,  250. 
Slag,  415,  541. 
Smalt,  552. 
Smaltite,  552. 
Smelling  salts,  438. 
Smelting,  414.     See  Metallurgy. 
Smithsonite,  482. 
Smokeless  powder,  328. 
Soap,  321,  323. 

And  hard  water,  473. 

Castile,  324. 

Hard,  323. 

Manufacture,  324. 

Soft,  323. 
Soda,  422,  423. 

Ash,  422. 

Baking,  422. 

Calcined,  422. 

Caustic,  424. 

Cooking,  423. 

Crystals,  424. 

Lime,  469. 

Mints,  424. 

Washing,  422. 


590 


INDEX 


Soda  water,  278. 
Sodium,  417. 

Acetate,  288,  319. 
Acid  carbonate,  422. 
Acid  sulphite,  344. 
Amalgam,  419,  425. 
Ammonium  phosphate,  396. 
And  water,  31,  61,  79. 
Arsenate,  403. 
Arsenite,  403. 
Bicarbonate,  281,  422,  423. 
Bromide,  365,  366. 
Carbonate,  417,  421,  422. 
Carbonate,  hydrolysis,  166,  167, 

422. 

Carbonate,  Leblanc  process,  421. 
Carbonate,  Solvay  process,  422. 
Chloride,  419,  420,  421,  426,  428, 

504. 
Chloride,  equilibrium  in  solution, 

193,  194. 
Chromate,  525. 
Cyanide,  330,  419,  427. 
Bichromate,  525. 
Dioxide,  427. 
Discovery,  417. 
Electrolytic    manufacture,    417, 

418. 

Hydride,  419. 
Hydroxide,  424,  425. 
Hydroxide,      electrolysis,      417, 

418. 

Hypochlorite,  209. 
Hyposulphite,  352. 
In  crust,  11,  417. 
lodate,  368,  370. 
Ions,  147,  425,  427. 
Manganate,  520. 
Manufacture,  417. 
Metaborate,  376. 
Metastannate,  507. 
Name,  417. 
Nitrate,  426,  427,  430. 
Nitrite,  427. 
Occurrence,  417. 
Organic  salts,  321,  323. 
Peroxide,  418,  419,  427. 
Plumbate,  515. 
Properties,  418. 
Silicate,  379,  383. 
Stanuate,  509, 


Sodium,  Sulphate,  421,  426. 

Sulphate,  electrolysis,  143. 

Sulphide,  421. 

Sulphite,  426. 

Sulphite,  acid,  344,  370. 

Test,  418. 

Thiosulphate,  352,  457. 

Tungstate,  527. 

Uranate,  527. 

Uses,  419. 
Soft  coal,  269. 

Water,  473. 
Solder,  508. 
Soldering,  376. 
Solubility,  denned,  67. 

Curve,  68. 

Degree,  63. 

Product,  195. 
Solute,  63. 
Solution,  63-76,  132-148,  193-195. 

And  chemical  action,  76,  184. 

And  crystallization,  70. 

And  dissociation,  163. 

And  electric  current,  132. 

And  electrolysis,  140. 

And  equilibrium,  189,  193-195. 

And  osmotic  pressure,  136. 

And  vapor  pressure,  73. 

Boiling  point,  140. 

Colloids,  382. 

Concentration  and  reaction,  185, 
193-195. 

Dissociation  table,  164. 

Electrolysis,  140. 

Electrolytic,  133. 

Freezing  point,  139,  239,  240. 

Gases,  64. 

General  properties,  132. 

Heat  of,  75. 

Liquids,  65. 

Non-electrolytic,  133. 

Normal,  163. 

Pressure,  452. 

Saturated,  67,  194,  195. 

Solids,  66-76. 

Summary,  147. 

Supersaturated,  69. 

Tension,  451. 

Terms,  63,  64. 

Thermal  phenomena,  75, 
Solvay  process,  422, 


INDEX 


591 


Solvent,  63. 

Spathic  iron  orc^548. 

Specific  gravity,  riietals,  412. 

Water,  56. 
Specific  heat,  245. 

And  atomic  weight,  246. 

Law,  245. 

Table,  246. 

Spectra,  inert  gases,  124. 
Spectroscope,  439. 
Spectrum,  439. 

Analysis,  438. 

Barium,  475. 

Calcium,  472. 

Line,  440. 

Potassium,  440. 

Radium,  476. 

Sodium,  440. 

Strontium,  474. 
Speculum  metal,  447. 
Spelter,  482. 
Sperrylite,  554. 
Sphalerite,  482. 
Spiegel  iron,  517. 
Spinels,  497. 
Spinthariscope,  529. 
Spontaneous  combustion,  24. 
Stahl,  25. 
Stalactite,  465. 
Stalagmite,  465. 
Standard  conditions,  42,  47. 
Stannic  compounds,  508,  509. 
Stannous  compounds,  490,  508,  509. 
Starch,  327. 

And  alcohol,  315. 

And  iodine,  327. 

Test,  370. 
Stas,  357. 
Stassfurt  deposits,  428. 

Bromine,  366. 

Calcium  chloride,  472. 

Chlorine,  196. 

Magnesium,  478. 

Minerals,  428. 

Potassium,  428. 

Sodium  sulphate,  426. 
Steam,  58. 
Stearin,  321,  322. 
Steel,  538. 

Acid  process,  541. 

Aluminium,  495, 


Steel,  Basic  process,  540,  541. 

Bessemer  process,  539. 

Cementation  process,  542. 

Composition,  538. 

Crucible  process,  542. 

Important  properties,  543. 

Open  hearth,  541. 

Siemens-Martin  process,  541. 

Special,  543. 

Tempering,  543. 

Thomas-Gilchrist  process,  540. 

Uses,  543. 
Sterling  silver,  455. 
Stibine,  405. 
Stibnite,  405,  406. 
Stoneware,  504. 
Storage  cell,  180. 
Stove  polish,  265. 
Strass,  388. 
Strontia,  474. 
Strontium,  473,  474,  475. 
Structural  formulas,  260,  312. 
Stucco,  471. 
Sublimate,  436. 

Corrosive,  489. 

Sublimation,     ammonium     chloride, 
436. 

Iodine,  370. 

Substitution,  32,  61,  103. 
Sucrose,  324. 
Sugar,  324. 

And  osmotic  pressure,  137. 

Beet,  325. 

Cane,  324. 

Fermentation,  315,  326. 

Fruit,  326. 

Granulated,  325. 

Grape,  326. 

Manufacture,  325. 

Molecular  weight,  241. 

Of  lead,  319. 

Refining,  325. 

Solution,  freezing  point,  139. 

Test,  326. 
Suint,  428. 

Sulphides,  332,  339,  340,  399,  437. 
Sulphites,  344. 

Ions,  344,  345. 
Sulphur,  332. 

Amorphous,  336. 

Chlorides,  335, 


592 


INDEX 


Sulphur,  Crystallized,  336. 

Dioxide,  340,  341,  342,  426. 

Extraction,  333. 

Flowers,  334. 

Forms,  336. 

Hexoxide,  343. 

Kiln,  333. 

Lime,  337. 

Milk  of,  337. 

Molecule,  335. 

Monoclinic,  336. 

Orthorhombic,  336. 

Phosphorus,  399. 

Properties,  2,  334. 

Purification,  334. 

Roll,  334. 

Spray,  337. 

Transition  point,  336. 

Trioxide,  343. 

Winning,  333. 

Sulphuretted  hydrogen,  337. 
Sulphuric  acid,  345. 

Calculation  formula,  100. 

Chamber  process,  346. 

Contact  process,  348. 

Di-,  352. 

Fuming,  351. 

Ions,  350,  351. 

Manufacture,  345. 

Nordhausen,  351. 

Plant,  346,  347,  348. 

Properties,  349. 

Pyro-,  352. 

Test,  351. 

Uses,  350. 

Sulphuric  ether,  316. 
Sulphurous  acid,  342. 

Ions,  344. 

Sun,  atmosphere,  30. 
Superphosphate  of  lime,  400. 
Supersaturation,  69. 
Sylvite,  428. 
Symbols,  10,  13,  97. 

And  ions,  134. 

Latin,  11. 

Tables,  10,  11,  12,  564. 
Synthesis,  water,  81. 

Table  salt,  419. 

Tables,  atomic  weights,  231,  564. 
Borax  beads,  376. 


Tables,  Calorific  value  of  coal,  269. 

Combination  by  volume,  225. 

Combining  water,  71. 

Composition  of  air,  118. 

Composition  of  body,  12. 

Composition  of  earth's  crust,  11. 

Composition  of  illuminating  gas, 
300. 

Composition  of  ocean,  11,  196. 

Dissociation,  164. 

Distribution  of  matter,  11. 

Equivalent   weights,    230,    231, 
232. 

Important  elements,  10. 

International     atomic     weights, 
564. 

lonization,  164. 

Ions,  147. 

Latin  symbols,  13. 

Less  common  elements,  12. 

Metalloids,  411. 

Metals,  411. 

Metric  equivalents,  559. 

Metric  system,  559. 

Metric  transformations,  559. 

Multiple  proportions,  91. 

Nitrogen  oxides,  221. 

Non-metals,  411. 

Periodic,  358. 

Solids  in  ocean,  53. 

Solubility  of  solids,  67. 

Specific  heat,  246. 

Valence,  elements,  254,  256. 

Valence,  radicals,  256. 

Vapor  pressure,  563. 

Water  in  food,  50. 
Tar,  coal,  295,  297,  329. 
Tartar,  319. 

Emetic,  320,  406. 
Tartaric  acid,  319. 
Tellurides,  458. 
Temperature,  critical,  128. 

Kindling,  25,  172. 

Low,  ethylene,  289. 

Standard,  41. 
Tempering,  543. 
Temporary  hardness,  473. 
Tension,  water  vapor,  58. 
Test,  21,  145,  208. 

Acetic  acid,  320. 

Alcohol,  320. 


INDEX 


593 


Test,  Aluminium,  502. 

Antimony,  406. 

Arsenic,  404. 

Barium,  475. 

Bismuth,  407. 

Borax  bead,  376. 

Boron,  375. 

Cadmium,  486. 

Calcium,  472. 

Carbon,  272. 

Carbon  dioxide,  276,  470. 

Chloride,  145. 

Chlorine  ions,  145. 

Chromium,  524,  525. 

Cobalt,  553. 

Copper,  450. 

Gold,  461. 

Hydrogen,  37. 

Hydrogen  sulphide,  340. 

Iodine,  327,  370. 

Iron,  549. 

Lead,  514,  515. 

Lithium,  435. 

Magnesium,  481. 

Manganese,  518,  520. 

Marsh's,  for  arsenic,  404. 

Mercury,  490. 

Nickel,  551. 

Nitrates,  220. 

Nitric  acid,  220. 

Oxygen,  21. 

Potassium,  429. 

Silicon,  384. 

Silver,  456. 

Sodium,  418. 

Starch,  327,  370. 

Strontium,  474. 

Sugar,  326. 

Sulphate  ions,  145. 

Sulphates,  351. 

Sulphuric  acid,  351. 

Tin,  509. 

Zinc,  485. 

Tetraboric  acid,  375. 
Tetrads,  257. 
Thenard,  87. 
Theory,  88. 

Atomic,  91. 

Electrolytic     dissociation, 

180. 
Thermal  equation,  176,  275. 


133, 


Thermal,  Phenomena  of  solution,  75. 
Thermit,  496. 
Thermometer,  560. 
Thomas-Gilchrist  process,  540. 

Slag,  541. 

Thorium,  515,  516,  531. 
Tin,  505. 

Alloys,  508. 

Amalgam,  508. 

Chlorides,  508,  509. 

Compounds,  508. 

Crystals,  508. 

Dioxide,  505,  506,  508. 

Disease,  506,  507. 

Displacement,  507. 

Foil,  508. 

Ions,  509. 

Metal  and  non-metal,  509. 

Mordants,  509. 

Oxymuriate,  509. 

Properties,  506. 

Salt  of,  508. 

Stone,  505. 

Stream,  508. 

Sulphides,  509. 

Test,  509. 

Transition  point,  506. 
Topaz,  497. 
Transition  point,  336. 

Sulphur,  336. 

Tin,  506. 

Transmutation,  458. 
Travertine,  465. 
Triacid  base,  158. 
Triads,  257. 
Tribasic  acid,  158. 
Tricalcium  phosphate,  391-394,  396, 

400. 

Trisilicate,  383. 
Trivalent  elements,  257. 
Tungsten,  527. 
Turnbull's  blue,  549. 
Turquoise,  497. 
Tuyeres,  534. 
Type  metal,  512. 

Univalent  elements,  254,  256,  257. 
Upward  displacement,  211. 
Uranium,  476,  527-531. 

Products,  531. 
Urea,  310. 


594 


INDEX 


Valence,  252,  253. 

And  atomic  weight,  254. 

And  equivalent  weight,  254. 

Determination,  253. 

Exceptions,  260. 

Radicals,  255. 

Raising  and  lowering,  545. 

Representation,  259,  312. 

Rules,  257,  258,  259. 

Tables,  254,  256. 

Variable,  256. 
Valentine,  Basil,  345. 
Vanadium,  543. 
Van  Helmont,  282. 
Vanillin,  316. 
Vapor  density,  236. 

And  molecular  weight,  236-238. 

Determination,  238,  239. 

Method,  236,  239. 
Vapor  pressure,  58-61. 

And  deliquescence,  74,  424. 

And  efflorescence,  72,  422. 

And  solution,  73. 

Table,  563. 

Vapors,  molecular  formula,  250. 
Vapor  tension,  58. 
Vaseline,  294. 

Velocity,  reaction,  184,  185,  188-191. 
Venetian  red,  546. 
Verdigris,  319,  450. 
Vermilion,  489. 
Victor  Meyer,  239. 
Vinegar,  317,  318. 
Vital  force,  310. 
Vitriol,  blue,  448. 

Green,  546. 

Oil  of,  345. 

White,  484. 
Volatile  alkali,  161. 
Volta,  176. 
Voltaic  cell,  178. 
Volume  equation,  252. 
Volumetric  composition,  81. 

Air,  118. 

Water,  81. 

Washing  soda,  422. 
Water,  50-85. 

Analysis,  54. 

And  calcium  oxide,  62. 

And  electric  current,  62,  143. 


Water,  And  ether,  66. 

And  hydrogen,  31,  36,  78. 

And  neutralization,  155. 

And  oxygen,  79. 

And  potassium,  429. 

And  sodium,  31,  61,  79. 

Baryta,  474. 

Bromine,  367. 

Catalysis,  186,  200. 

Chalybeate,  53. 

Chemical  properties,  61. 

Chlorine,  198. 

Composition,  78-85. 

Crystallization,  71,  98. 

Decomposition,  61,  80,  419. 

Density,  56. 

Dissociation,  165,  194,  422.  448, 

449. 

Distillation,  54. 
Drinking,  53. 
Electrolysis,  143. 
Evaporation,  57-61. 
Expansion,  56. 
Gas,  283,  298,  299,  300. 
General  properties,  50-76. 
Glass,  383. 

Gravimetric  composition,  83. 
Hard,  281. 
Hardness,  473. 
Hydrogen  sulphide,  338. 
Hydrolysis,  166. 
Industrial  application,  51. 
In  food,  50. 
In  nature,  50,  51. 
Ions,  165,  194,  422,  448,  449. 
Javelle's,  209. 
Lime,  469. 
Lithia,  435. 
Mineral,  52. 
Natural,  52. 
Ocean,  11,  53. 
Of  crystallization,  71,  98. 
Pure,  properties,  56. 
Purification,  499. 
River,  53. 
Royal,  224. 
Soda,  278. 
Soft,  473. 
Solvent  power,  63. 
Underground,  52,  281,  465. 
Vapor  and  catalysis,  86. 


INDEX 


595 


Water,  Vapor  in  atmosphere,  120. 

Vapor,  molecular  equation,  251, 
252. 

Vapor.     See  Vapor  pressure. 

Volumetric  composition,  81. 
Watt,  83, 
Weight,  liter  of  air,  115. 

Ammonia  gas,  211. 

Carbon  dioxide,  277. 

Chlorine,  198. 

Formula,  238. 

Hydrogen,  33. 

Hydrogen  chloride,  205. 

Hydrogen  sulphide,  338. 

Molar,  238. 

Oxygen,  19,  48,  59. 

Nitrogen,  129. 

Sulphur  dioxide,  342. 
Weldon  process,  197,  518. 
Welsbach  light,  309,  516. 
Whisky,  315. 
White  arsenic,  402. 

Lead,  513,  514. 

Metal,  508. 

Paint,  513. 

Vitriol,  484. 
Whitewash,  470. 
Whiting,  467. 
Willemite,  476,  482. 
Willson,  285. 
Wine,  315. 
Witherite,  474. 
Wohler,  310. 
Wollastonite,  383. 
Wood  alcohol,  314. 

Ashes,  428,  432. 


Wood,  Charcoal,  271,  272,  273. 

Distillation,  272,  317. 

Petrified,  379. 

Silicified,  379. 

Vinegar,  317. 
Wood's  metal,  486, 
Wrought  iron,  537, 

X-rays,  529. 
Xenon,  123,  358. 

Yeast,  315,  327. 

Zero,  absolute,  44,  560. 
Zinc,  482. 

Alloys,  484. 

Blende,  482. 

Chloride,  484. 

Determination    atomic    weight, 
247. 

Displacement,  451,  483. 

Dust,  482,  484. 

Family,  491. 

Hydroxide,  485. 

Ions,  451,  485. 

Oxide,  482,  483,  484. 

Properties,  482,  485. 

Silicate,  482. 

Sulphate,  451,  484. 

Sulphide,  476,  482,  484,  529. 

Test,  485. 

White,  484. 
Zincates,  483,  485. 
Zincite,  482. 
Zircon,  383. 
Zymase,  315. 


sS^g^Mgwass 


THE  UNIVERSITY  OF  CALIFORNIA  LIBRARY 


